2: Bonds, James Bonds How can we apply the physical and chemical properties of elements to predict the formation of compounds?

Complete the physical versus chemical properties hand out with the person sitting next to you.
Drill 1 10/14 (A Day) 10/16 (B Day) Outcome: I can explain the properties of the elements using a periodic table. Goal: CW 1, CW 2

CW 1: Mendeleev’s Table Video Link: Was Mendeleev the only person to sort the elements into lists by similar properties? What is the significance of the dashes Mendeleev placed on his periodic table? How could Mendeleev predict the properties of eka-aluminum? What is the relationship between gallium and eka-aluminum?

CW 2: Properties of the Periodic Table
Complete the assigned questions in the unit packet by read along on the page. Go to LeffelLabs/Unit 2 and click on the PDF for CW 2.

Delocalized bonding: electrons involved in bonding are shared by all atoms, nuclei are little islands held together by an electron sea. Metals conduct because these electrons can flow freely Malleable/ ductile because the atoms aren’t rigidly locked in place and can shift without destroying the network of bonds Shiny because electrons absorb and emit light energy Localized bonding: electrons are held in one place between two atoms Poor conductors because electrons can’t flow Brittle because shifting around these atoms would break bonds between electrons Metalloids are in between localized and delocalized bonding So the properties are in between metals and nonmetals Semiconductors: conduct somewhat, not as much a metals, computing applications

Journal Write 1 Create a table to compare the type of bonds in metalloids to those in metals (delocalized) and nonmetals (localized).

Summary 1 10/14 (A Day) 10/16 (B Day)
Complete CW 1 and CW 2. Who Tagged the Lab Bench Conclusion, 10/19 (B Day) SciResearch: 6 Develop a Procedure, 10/25 (B Day) SciResearch: 7 Get Approval to Start, 10/31 (B Day) Summary 1 10/14 (A Day) 10/16 (B Day) Outcome: I can explain the properties of the elements using a periodic table. Goal: CW 1, CW 2

Drill 2 10/17 (A Day) 10/18 (B Day) 15 periwinkles = 5 snails
12 snails = 1 fish 300 fish = 1 shark How many sharks are in periwinkles? How many snails are in 3 sharks? Drill 2 10/17 (A Day) 10/18 (B Day) Outcome: I can apply electron configurations to explain periodic trends. Goal: CW 3, CW 4 Hand In: Who Tagged the Lab Bench Conclusion 2500 𝑝𝑒𝑟. 1 × 5 𝑠𝑛𝑎𝑖𝑙𝑠 15 𝑝𝑒𝑟. × 1 𝑓𝑖𝑠ℎ 12 𝑠𝑛𝑎𝑖𝑙𝑠 × 1 𝑠ℎ𝑎𝑟𝑘 300 𝑓𝑖𝑠ℎ = 𝑠ℎ𝑎𝑟𝑘𝑠 3 𝑠ℎ𝑎𝑟𝑘𝑠 1 × 300 𝑓𝑖𝑠ℎ 1 𝑠ℎ𝑎𝑟𝑘 × 12 𝑠𝑛𝑎𝑖𝑙𝑠 1 𝑓𝑖𝑠ℎ =10800 𝑠𝑛𝑎𝑖𝑙𝑠

CW 3: Electron Configuration
Explain the atom according to the quantum mechanical model. Updates the Bohr model: electrons are NOT in circular orbits around the nucleus Describes the location of electrons in atoms based on probability What is an atomic orbital? 3D cloud of moving electrons around the nucleus The electron cloud is denser where electrons are more likely to be found, a fast blur of electrons Each atomic orbital differs in amount of energy and shapes (S, P, D, F), orbitals “stack”

How does the electron configuration relate to the arrangement of electrons in atoms? Shows how e– are distributed among energy levels and atomic orbitals # of e- in energy level Energy Level 1s2 Shape/ name of orbital

S P D F

What is the electron configuration of the following? Magnesium Iron Krypton Rubidium

Ca: Number of valence electrons: 1s2 2s2 2p6 3s2 3p6 4s2 Ca2+: Number of valence electrons: 1s2 2s2 2p6 3s2 3p6 4s2 2 8

S: Number of valence electrons: 1s2 2s2 2p6 3s2 3p4 S2–: Number of valence electrons: 1s2 2s2 2p6 3s2 3p6 6 8

CW 4: Trends in the Periodic Table
Take notes in the table below. Property Definition Trend Across a Row (left to right) Trend Down a Column Atomic Radius Radius of an atom from the nucleus to the edge of the electron cloud Decreases Increases Ionization Energy Minimum amount of energy required to remove the most loosely bound (valence) electron to form a cation. Increases Electronegativity How well an atom attracts electrons in a bond Metallic Character How easily an atom loses electrons (metals react by losing electrons) Decreases

Atomic Radius

Ionization Energy

Electronegativity

For each of the following, list the element which has the highest: Atomic Radius: F, Cl, Br Ionization Energy: K, Ca, Sc Electronegativity: C, Si, P, N

What is shielding? E- in lower energy levels “shield” the outer e- from the nucleus Down a column = more energy levels = more shielding = less attraction to nucleus

Li: 1s2 2s1 Na: 1s2 2s2 2p6 3s1 K: 1s2 2s2 2p6 3s2 3p6 4s1 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? How does shielding affect the magnitude of attraction between the nucleus and the valence electrons? If the attraction between nucleus and valence electron is weak, would that electron be found closer to the nucleus, or farther away? Explain.

How does shielding affect the following? Atomic Radius: more layers = weaker attraction = larger atom Ionization Energy: more layers = weaker attraction = lower ionization energy Electronegativity: more layers = weaker attraction = lower electronegativity

What is nuclear charge? Charge of nucleus; adding more p+ (across a row) = more positive nucleus = more attraction between e- and nucleus +3 +9

Consider the atoms above. Which would have a larger positive nuclear charge? How does nuclear charge affect the magnitude of attraction between the nucleus and the valence electrons? If the attraction between nucleus and valence electron is strong, would that electron be found closer to the nucleus, or farther away? Explain. Atom K Ca Sc Ti V # of p+

How does nuclear charge the following? Atomic Radius: more protons = stronger attraction = electrons move in = smaller radius Ionization Energy: more protons = stronger attraction = higher ionization energy Electronegativity: more protons = stronger attraction = higher electronegativity

Journal Write 2 Complete the summary. How does increasing shielding affect the attraction between the nucleus and electrons? How does increasing nuclear charge affect the attraction between the nucleus and electrons?

HW 1: Electron Configuration Complete CW 3 and CW 4. Who Tagged the Lab Bench Conclusion, 10/19 (B Day) SciResearch: 6 Develop a Procedure, 10/25 (B Day) SciResearch: 7 Get Approval to Start, 10/31 (B Day) Summary 2 10/17 (A Day) 10/18 (B Day) Outcome: I can apply electron configurations to explain periodic trends. Goal: CW 3, CW 4 Hand In: Who Tagged the Lab Bench Conclusion

What are the electron configurations of the elements below?
K Na F Cl Drill 3 10/21 (A Day) 10/22 (B Day) 1s2 2s2 2p6 3s2 3p6 4s1 1s2 2s2 2p6 3s1 Outcome: I can use electron configuration to explain periodic trends and predict bond types. Goal: CW 4, CW 5 Hand In: SciResearch Page 6 1s2 2s2 2p5 1s2 2s2 2p6 3s2 3p5

CW 5: What Kind of Bond will Form?
Define electronegativity How strongly an atom attracts electrons in a chemical bond. Pauling Electronegativity Scale

Use the information below and the electronegativity value table on the previous page to complete the chart below. First Element Metal or Nonmetal? 1st Element Electroneg. Second Element 2nd Element Electroneg. Electroneg. Difference Bond Type Ca Metal 1.0 F Nonmetal 4.0 3.0 Ionic Ag 1.9 Au 2.4 0.5 Metallic H 2.1 Cl 0.9 Polar Covalent Cu 0.5 Nonmetal 2.1 Nonpolar Covalent Mg 1.2 1.8 O 3.5 1.4 Very Polar Covalent

Describe the electronegativity difference and the type of element (metals or nonmetals) involved in each of the types of bonds. Bond Type Electronegativity Difference Types of Elements Involved Ionic Polar Covalent Nonpolar Covalent Metallic Very Large, ≥ 2.0 Metal and Nonmetal Moderate, 0.4 to 2.0 Nonmetals Very small, 0.0 to 0.4 Nonmetals Not important! Metals

Electronegativity Difference δ+ δ − +1 −1 0.4 2.0 nonpolar covalent bond polar covalent bond ionic bond larger distribution of charge bonds become more “ionic like”

Use the periodic table below. Find and highlight Na and Cl in yellow. Find and highlight C and O in green. Which pair would form an ionic bond? Na (metal) and Cl (nonmetal) Which pair would form a covalent bond? C (nonmetal) and O (nonmetal) Considering the position of the elements you highlighted, which pair of elements would you expect to have a large electronegativity difference? Explain. The farther apart, the greater the electronegativity difference: the metal (Na) and the nonmetal (Cl)

Complete the table below. How many valence electrons are needed for a stable electron configuration? A full octet: 8 valence electrons Fluorine has 7 valence electrons. How does this explain the very high electronegativity value of fluorine? It is close to a full octet, wants to gain an electron, not lose any. How many valence electrons would sodium have if it lost its 3s1 electron, forming the cation Na1+? Is this configuration more stable, or less stable? If the 3s1 e– was lost, the highest energy level would be 2, which has a full octet, making it more stable. Element Electron Configuration # Valence Electrons Electronegativity Value (Page 7) Na 1s2 2s2 2p6 3s1 1 0.9 F 1s2 2s2 2p5 7 4.0

Journal Write 3 Considering your answer to Question 8, why does sodium have such a low electronegativity value?

CW 6: Bonding Types Virtual Lab
Go to LEFFELlabs/ unit 2 to access the virtual lab. In the table, you should summarize and explain each property. Your goal is to understand, not memorize. This is on the quarterly assessment.

Review the formation of each bond type. Summarize below. Include a labeled picture for each summary. Ionic One atoms loses electrons, another gains those electrons. Leads to cations and anions, opposites attract, forms ionic bond. Metallic Each metal atom loses electrons and forms a cation, electrons are shared amongst all metal cations in an electron sea

Review the formation of each bond type. Summarize below. Include a labeled picture for each summary. Covalent (explain the difference in electron sharing for polar and nonpolar) Polar Electrons are shared between atoms unequally due to different electronegativities. Nonpolar Electrons are shared between atoms equally due to similar electronegativities.

Explain the relationship between covalent bond and molecular compound. A covalent bond occurs between two nonmetal atoms. A molecular compound is held together by covalent bonds.

Property Ionic Metallic Covalent: Polar Covalent: Nonpolar Dissolves in a polar solvent (water) Yes: polar water interacts with ionic charges No: metals do not dissolve Yes (like dissolves like) No (like dissolves like) Dissolves in a nonpolar solvent (oil) No: ionic charges cannot interact with uncharged solvents Conduct in the solid state No: ions must move to conduct; not possible as a solid Yes: electrons move in a “sea”, able to flow and conduct No: no moving ions/ charges Conducts in the aqueous state Yes: ions dissolve and move in aqueous solutions, can conduct No: doesn’t dissolve in water No: doesn’t dissolve, no ions/ charges Melting point/ Boiling point High: very strong attractions between ions High: strong mutual attraction between electron sea and metal cations Low: weaker attractions hold these compounds together Malleability No: ions may repel and shatter the solid if like charges touch Yes: electron sea allows cations to slide past each other Solids will be brittle, liquids and gases don’t have malleability

Complete CW 4 and CW 5 HW: Complete CW 6 SciResearch: 6 Develop a Procedure, 10/25 (B Day) SciResearch: 7 Get Approval to Start, 10/31 (B Day) Summary 3 10/21 (A Day) 10/22 (B Day) Outcome: I can use electron configuration to explain periodic trends and predict bond types. Goal: CW 4, CW 5 Hand In: SciResearch Page 6

Ms. L’s cat has food allergies, and can no longer eat fish
Ms. L’s cat has food allergies, and can no longer eat fish. Abby normally eats 1.3 cans of tuna per day. 1/2 can of tuna = 26 calories 1 oz. of rabbit meat = 58 calories How much rabbit meat does Abby need to eat per week? Drill 4 10/23 (A Day) 10/24 (B Day) Outcome: I can illustrate the formation of a covalent bond. Goal: CW 7, CW 8 1.3 𝑐𝑎𝑛𝑠 𝑡𝑢𝑛𝑎 1 𝑑𝑎𝑦 × 26 𝑐𝑎𝑙. 0.5 𝑐𝑎𝑛𝑠 𝑡𝑢𝑛𝑎 × 1 𝑜𝑧. 𝑟𝑎𝑏𝑏𝑖𝑡 58 𝑐𝑎𝑙. × 7 𝑑𝑎𝑦𝑠 1 𝑤𝑒𝑒𝑘 = = 𝑜𝑧. 𝑟𝑎𝑏𝑏𝑖𝑡 𝑤𝑒𝑒𝑘 =8.2 𝑜𝑧. 𝑟𝑎𝑏𝑏𝑖𝑡 𝑤𝑒𝑒𝑘

Flame Test Lab Feedback
Purpose = one direct sentence Font size 12 or 14 Underline materials: there is no separate materials list Put your headings in the same order as the rubric Include last names of partners Answer each conclusion question separately References go AFTER the conclusion Save files in an organized way. “Lab Report” is a terrible file name

CW 7: Covalent Bonds Gizmo
Go to Username: leffel Password: leffel Complete the gizmo: READ!

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? Think and discuss: How many electrons are represented by the line between each atom? What do the other dots around each atom represent? One line = 2 electrons Other dots = electrons not being shared F F F F

Observe: Like fluorine and most other elements, oxygen atoms are most stable with a full complement of eight valence electrons. How many valence electrons does each oxygen atom have now? How many more electrons does each oxygen atom need to be stable? 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? 6 2 2 pairs 2 lines

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. O O O O O O

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. Element Symbol # of valence electrons Valence Sum Fluorine F Hydrogen H Oxygen O Nitrogen N Chlorine Cl Carbon C Silicon Si 7 1 8 1 1 2 6 2 8 5 3 8 7 1 8 4 4 8 4 4 8

Practice: Create covalent bonds and stable molecules for the remaining substances. Draw Lewis diagrams for each one.

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? These gases already have a full valence (full octet) and do not need to share electrons to become stable.

CW 8: Naming Molecular Compounds
Based on the examples above, what kinds of atoms form molecular compounds (Metals? Nonmetals? Both?) Nonmetals combined with nonmetals. What do the following prefixes mean? 1 4 Mono Tetra Di Penta Tri Hexa 2 5 3 6

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 CCl4 H2S Pl3 SBr2 BF3 N2 P4O10 SF6

Name the molecular compounds below using the rules above. Formula Name N2H4 Dinitrogen tetrahydride SBr2 Sulfur dibromide XeF4 Xenon tetrafluoride P4O3 Tetraphosphorus trioxide NH3 Nitrogen trihydride BCl3 Boron trichloride CO2 Carbon dioxide H2O Dihydrogen monoxide

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

HW 2: Molecular Compounds Complete CW 7 and CW 8 SciResearch: 7 Get Approval to Start, 10/30 (B Day) Summary 4 10/23 (A Day) 10/24 (B Day) Outcome: I can illustrate the formation of a covalent bond. Goal: CW 7, CW 8

Drill 5 10/25 (A Day) 10/28 (B Day) Mg Ca F Cl
Write the electron configuration for the following: Mg Ca F Cl Now, draw the Lewis dot structure for each atom. Drill 5 10/25 (A Day) 10/28 (B Day) Mg 1s2 2s2 2p6 3s2 Outcome: I can predict the 3D shape of molecules using a Lewis diagram. Goal: CW 9 Ca 1s2 2s2 2p6 3s2 3p6 4s2 F 1s2 2s2 2p5 1s2 2s2 2p6 3s2 3p5 Cl

CW 9: Lewis Dot Structures (Molecular)
Write the number of valence electrons above each column.

How many electrons are shared in a single bond? 2 electrons How many electrons are shared in a double bond? 4 electrons How many electrons are shared in a triple bond? 6 electrons What type of bond is always formed with hydrogen? Single bond, hydrogen has one electron and needs one more to be stable What is the main difference between drawing Lewis structures for ionic compounds and covalent molecules? In ionic compounds, atoms have charges due to the transfer of electrons. In covalent compounds, electrons are shared, not transferred, so there are no charges.

Molecule CO2 HCN CH2O Draw the Lewis dot structure of the individual atoms. Find the number of bonds each atom can form (each unpaired electron = 1 bond). Determine the central atom (atom with the fewest electrons, never H). Draw the Lewis structure (place electrons in bonds, then place remaining electrons around their correct atoms).

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

Compare the 3D model you built to the shapes shown on the VSEPR chart. Do they match up? More or less – but the models are missing non bonding electrons. What is missing from the models that you can see in the Lewis dot structure? The Lewis dot structure shows the lone pairs, while the models do not.

HW 3: IMF Notes (graded!) Complete CW 9 Journals 1 to 4 collected next class For journal 4, write a summary of the rules. SciResearch: 7 Get Approval to Start, 10/30 (B Day) Summary 5 10/25 (A Day) 10/28 (B Day) Outcome: I can predict the 3D shape of molecules using a Lewis diagram. Goal: CW 9

Draw the Lewis diagram for C2H4 (hint: 2 “central” atoms).
What is the name of this compound? Drill 6 10/29 (A Day) 10/30 (B Day) Outcome: I can explain the properties of substances using intermolecular forces. Goal: CW 10 Hand In: SciResearch Page 7

Exit Tickets Do not lose this packet as we will be using it throughout 1st and 2nd marking period. Complete exit tickets when you are told to do so. You will need the given conversion factors, as well as a table of elements. You must use dimensional analysis to receive credit. All grades are based on correctness. Grades may count as assessment or classwork based on teacher discretion. Pop quizzes on dimensional analysis are possible and likely. A major test concerning dimensional analysis will be given before semester exams.

Harvest for the Hungry What if you had to choose?
Buy food or buy gasoline to get to work? Buy food or go to the doctor? Buy food or pay rent? For a family of four (two adults, two children), both adults have to work 81 hours per week at minimum wage in the state of Maryland to make a living wage. For a family of five (two adults, three children), both adults must work 93 hours to make a living wage

Harvest for the Hungry All students will complete the dimensional analysis questions for a grade, based on correctness. Due date: 11/13 (A day) or 11/14 (B day). Regardless of grade or service learning hour status, the calculations will count as an assessment grade. Any student found copying this assignment will receive a grade of zero; shall not be able to redo this assignment; and will be reported for academic dishonesty. Sophomores must do the following in order to earn credit for service learning. Fundraise for as much money as possible or collect as many canned goods as possible (at least enough to feed a meal to a family of six) You must calculate (using dimensional analysis and showing all work) how many meals you are providing. Put your calculations on a sheet of paper and put it in your bag of food with your name and period if you want credit. Due date: 11/13 (A day) or 11/14 (B day). If you fail to complete the requirements your name will be submitted to the office as failing service learning. This is a requirement for graduation. Items will not be accepted after 11/16. If you cannot collect money or canned goods, you will need to complete 10 hours of volunteer work, complete the forms (including having a supervisor sign them), and hand them in to your teacher by the end of the semester (1/26). Class that brings the greatest total amount = donuts & coffee

HW 3: Intermolecular Forces Notes
What are intramolecular forces? Attractive forces that within a molecule (chemical bonds) What are intermolecular forces? Attractive forces between two different molecules Explain how a molecule can have a permanent dipole. One atom is more electronegative than the other, attracts electrons = partial negative charge. The other atom is electron poor = partial positive charge.

What is a dipole-dipole interaction and how does it form? Partial positive pole of one molecule attracts the partial negative pole of another molecule. δ+ δ- δ+ δ-

Explain hydrogen bonding. How is it similar to a dipole-dipole interaction? Partial positive is H, partial negative is F, O, or N. Attraction is stronger than a typical dipole-dipole, so it gets a special name: hydrogen bonding. δ+ F, O, or N δ- H δ+ F, O, or N δ- H

What is a nonpolar molecule? Molecule with no dipole – no distribution of charge How can a molecule be nonpolar due to similar electronegativities? Similar electronegativities = no atom wins the “tug of war” over the shared electrons in the bond

How can a molecule be nonpolar due to symmetry? Polar bonds pull the electrons in opposite directions, cancelling each other out

How do dispersion forces occur in molecules? Electrons are swarming around each atom - might “catch” the atom with an uneven distribution of electrons Creates a short lived, temporary dipole = short lived temporary attractions Cl δ+ δ- Cl δ+ δ-

Explain why water is a liquid and not a gas. Water has all 3 IMFs, holding it together as a liquid Why does water have such a high boiling point compared to similar molecules? Having all 3 IMFs = hard to separate water molecules far enough to become a gas = need lots of heat

CW 10: Inter- and Intra- Molecular Forces

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. When liquid water evaporates into gaseous water, are any bonds between H atoms and O atoms within a molecule broken? No, it’s still H2O On average, are the intermolecular forces stronger in liquid water or gaseous water? Liquid, because the molecules are closer together and can interact more

Consider the two molecules below to complete the table. Molecule Picture Intramolecular Forces Intermolecular Forces Water H2O Polar Covalent Nonpolar Covalent Ionic Hydrogen bonding Dipole-Dipole Dispersion Hydrogen sulfide H2S

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. Weaker attractions = gas H2S cannot hydrogen bond (like H2O), meaning it has weaker attractions

Methane and methanol have nearly the same chemical formula. Identify the forces present in each molecule. In terms of intermolecular forces, why is methane a gas and methanol is a liquid? Methanol can hydrogen bond, these stronger attractions hold the methanol molecules close together as a liquid. Molecule Picture Intramolecular Forces Intermolecular Forces Methane CH4 Polar Covalent Nonpolar Covalent Ionic Hydrogen bonding Dipole-Dipole Dispersion Methanol CH3OH

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. 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? H2O forms a bead due to the strong attractions between the water molecules. Isopropyl alcohol spreads out, meaning the isopropyl alcohol molecules are less attracted to each other than in water.

IMFs and Phase Changes Evaporization occurs when the attractive forces between molecules are weakened enough that molecules are able to escape the liquid phase and enter the gas phase. Using your finger “push up” the two drops from Question 7 about two inches, so you have two “streaks”. Observe which substance evaporates more quickly. Isopropyl alcohol evaporates faster In Question 7, which substance had the stronger intermolecular forces? Does this agree with your observations about which evaporated more quickly? Isopropyl alcohol had weaker IMFs. This agrees; isopropyl alcohol evaporated first, as there are weaker attractions to hold the molecules together as a liquid.

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

Which molecule is “better” at forming hydrogen bonds? Each water molecule can form 4 hydrogen bonds, while isopropyl alcohol molecule can only make 2 hydrogen bonds. How does this explain the difference in evaporization time between water and isopropyl alcohol? Even though both form hydrogen bonds, water does it better, so it will take more time to evaporate. 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. 65°C: Methanol has all 3 IMFs, so it takes a lot of heat to weaken them to the point of boiling. -161.5°C: Methane only has dispersion forces, so it is easier to boil.

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 a large number of 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.

Waft the following substances to determine if they are volatile or not. Then determine which IMFs are present in each molecule.

Journal Write 5 Explain your observations of the volatility of the substances in terms of the intermolecular forces present in each molecule.

HW 4: Review for the Quarterly Assessment Video on LEFFELlabs Due 11/5 Complete CW 10 Complete Journals 1 – 3 and 5 Collected next class Skip journal 4 SciResearch: 7 Get Approval to Start, 10/30 (B Day) Summary 6 10/29 (A Day) 10/30 (B Day) Outcome: I can explain the properties of substances using intermolecular forces. Goal: CW 10 Hand In: SciResearch Page 7

Agenda 10/31 (A) and 11/1 (B) Review

Agenda 11/4 (A) and 11/5 (B) Quarterly Assessment

Drill 7 11/6 (A Day) 11/7 (B Day) What is the e− configuration for Cl?
What is the short hand e− configuration for Cl? How many e− are in the valence energy level? How many valence e− are needed for an atom to be stable? Drill 7 11/6 (A Day) 11/7 (B Day) Outcome: I can predict the formula of an ionic compound. Goal: CW 11, CW 12

CW 11: Lewis Dot Structures (Ionic)
Element Property Before Making an Octet After Making an Octet Na Electron configuration 1s2 2s2 2p6 3s1 1s2 2s2 2p6 # Protons 11 # Electrons 10 Charge +1 Lewis Dot Structure Na

Element Property Before Making an Octet After Making an Octet Mg Electron configuration 1s2 2s2 2p6 3s2 1s2 2s2 2p6 # Protons 12 # Electrons 10 Charge +2 Lewis Dot Structure Mg

Element Property Before Making an Octet After Making an Octet Al Electron configuration 1s2 2s2 2p6 3s2 3p1 1s2 2s2 2p6 # Protons 13 # Electrons 10 Charge +3 Lewis Dot Structure Al

Explain why these atoms tend to form cations. Their valence shells are somewhat empty (3 e– or fewer), so they must lose valence electrons to get a full octet Losing electrons = positive (cation)

Element Property Before Making an Octet After Making an Octet N Electron configuration 1s2 2s2 2p3 1s2 2s2 2p6 # Protons 7 # Electrons 10 Charge 3– Lewis Dot Structure N

Element Property Before Making an Octet After Making an Octet O Electron configuration 1s2 2s2 2p4 1s2 2s2 2p6 # Protons 8 # Electrons 10 Charge 2– Lewis Dot Structure O

Element Property Before Making an Octet After Making an Octet F Electron configuration 1s2 2s2 2p5 1s2 2s2 2p6 # Protons 9 # Electrons 10 Charge 1– Lewis Dot Structure F

Explain why these atoms tend to form anions. Their valence shells are somewhat full (5 e– or more), so they must gain valence electrons to get a full octet Gaining electrons = negative (anion) Would neon form an anion? Explain why or why not. No, neon is a noble gas with 8 electrons Neon doesn’t need to gain or lose electrons

Draw the Lewis dot structure for the following ionic compounds. NaBr MgO +1 -1 Na Na: 1s2 2s2 2p6 3s1 [ ] Na Br Br: 1s2 2s2 2p6 3s2 3p6 4s2 4p5 Br Mg +2 -2 Mg: 1s2 2s2 2p6 3s2 [ ] Mg O O: 1s2 2s2 2p4 O

Draw the Lewis dot structure for the following ionic compounds. Na3N AlF3 Na -3 +1 [ ] +1 Na: 1s2 2s2 2p4 3s1 Na N Na N N: 1s2 2s2 2p3 +1 Na Na Na: 1s2 2s2 2p4 3s1 Na Na: 1s2 2s2 2p4 3s1 Al: 1s2 2s2 2p6 3s2 3p1 F -1 -1 [ ] +3 [ ] F: 1s2 2s2 2p5 F Al F Al F F: 1s2 2s2 2p5 [ ] F F: 1s2 2s2 2p5 F -1

Ionic compounds form a crystal lattice structure of ions.

CW 12: Build an Ionic Formula
Polyatomic ions: a group of covalently bonded atoms that tend to stay together as if they were a single ion.

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 PO4-3 3– Sulfate SO4-2 S:1 O:4 2– Acetate C2H3O2 C:2 H:3 O:2 Carbonate CO3-2 C:1 O:3

The shortcut to finding the valence electrons! Li: 1s2 2s1 Na: 1s2 2s2 2p6 3s1 K: 1s2 2s2 2p6 3s2 3p6 4s1 Be: 1s2 2s2 Mg: 1s2 2s2 2p6 3s2 Ca: 1s2 2s2 2p6 3s2 3p6 4s2 B: 1s2 2s2 2p1 C: 1s2 2s2 2p2 N: 1s2 2s2 2p3 O: 1s2 2s2 2p4 F: 1s2 2s2 2p5 Ne: 1s2 2s2 2p6 All have 1 valence electron All have 2 valence electrons 3 valence electrons 4 valence electrons 5 valence electrons 6 valence electrons 7 valence electrons 8 valence electrons

+1 1 8 X +2 2 3 +3 ? 4 5 -3 -2 6 7 -1 Ignore these – more complex e- configuration – no easy way to predict the valence.

Use the size of the Lego to determine the charge. Cations are positive because they LOSE electrons (metals). Anions are negative because they GAIN electrons (nonmetals). When we filled out the table of elements with ion charges, we skipped the transition metals. Transitions metals may have more then one possible charge.

Use the following blocks to represent your ions. Indicate the charge of each. Cations Formula and Charge Anions Light pink – calcium Ca2+ Black – sulfate SO42– Lime – sodium Na+ Gray – chloride Cl– Light Blue – copper Cu+ Cu2+ Orange – acetate C2H3O2– Dark Blue – aluminum Al3+ Yellow – oxide O2– Green – lead Pb4+ Red – nitride N3–

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 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 Al+3 Cl-1 1 3 +3 -3 AlCl3 Al+3 Cl-1 Al+3 Cl-1 1 3 Cl-1 AlCl3 Cl-1

Combining Ions Cation Formula Anion Formula Number of Cations Number of Anions Total + charge Total – charge Ionic Formula Aluminum chloride Al+3 Cl-1 1 3 +3 -3 AlCl3 Calcium sulfate Ca+2 SO4-2 +2 -2 CaSO4 Lead (IV) oxide Pb+4 O-2 2 +4 -4 PbO2 Copper (II) chloride Cu+2 CuCl2 Aluminum oxide +6 -6 Al2O3 Copper (II) nitride N-3 Cu3N2 Sodium sulfate Na+1 Na2SO4 Calcium nitride Ca3N2

CW 12: Build an Ionic Compound
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?

Complete CW 11 and CW 12 HW 5: Ionic Compounds Summary 7 11/6 (A Day) 11/7 (B Day) Outcome: I can predict the formula of an ionic compound. Goal: CW 11, CW 12

The average American student is in class 330 minutes/day
The average American student is in class 330 minutes/day. How many hours/year is this? (1 school year = 180 days) Drill 8 11/8 (A Day) 11/11 (B Day) Outcome: I can write names and chemical formulas for ionic compounds Goal: CW 13

CW 13: Naming Ionic Compounds
Explain why sodium obtains a stable electron configuration by losing its valence electron. By losing the 3s1 electron, the next energy level (2s2 2p6) has a full octet. Na: 1s2 2s2 2p6 3s1

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? Li N Be O B F C Cl L 1 +1 G 3 -3 L 2 +2 G 2 -2 L 3 +3 G 1 -1 ? 4 ? G 1 -1

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. +1 +2 +3 -3 -2 -1

Why is it difficult to assign an oxidation state to group 14? That column has 4 valence electrons – we are not sure if it will lose 4 or gain 4 electrons.

Name to Formula Write the cation and the anion. Watch out for polyatomic ions. Determine the charges using romans OR periodic table. Criss Cross Apple Sauce. Magnesium Acetate +2 -1 Mg (C2H3O2) 1 2 Mg(C2H3O2)2

Name to Formula Write the cation and the anion. Watch out for polyatomic ions. Determine the charges using romans OR periodic table. Criss Cross Apple Sauce. Calcium Phosphate +2 -3 Ca (PO4) 3 2 Ca3(PO4)2

Name to Formula Write the cation and the anion. Watch out for polyatomic ions. Determine the charges using romans OR periodic table. Criss Cross Apple Sauce. Aluminum Sulfide +3 -2 Al S 2 3 Al2S3

Name to Formula Write the cation and the anion. Watch out for polyatomic ions. Determine the charges using romans OR periodic table. Criss Cross Apple Sauce. Iron (III) Oxide +3 -2 Fe O 2 3 Fe2O3

Name to Formula Write the cation and the anion. Watch out for polyatomic ions. Determine the charges using romans OR periodic table. Criss Cross Apple Sauce. Lead (II) Sulfate +2 -2 Pb (SO4) 2 2 PbSO4

Name to Formula Write the cation and the anion. Watch out for polyatomic ions. Determine the charges using romans OR periodic table. Criss Cross Apple Sauce. Vanadium (V) Phosphate +5 -3 V (PO4) 3 5 V3(PO4)5

Name to Formula Write the cation and the anion. Watch out for polyatomic ions. Determine the charges using romans OR periodic table. Criss Cross Apple Sauce. Copper (II) Chloride +2 -1 Cu Cl 1 2 CuCl2

Formula to Name Write the name of the cation If the cation is a transition metal, determine the charge (write in romans) Write the name of the anion, changing the “-ine” to “-ide” Watch out for polyatomic ions (just write them) NaCl Sodium Chlorine ide Sodium Chloride

Formula to Name Write the name of the cation If the cation is a transition metal, determine the charge (write in romans) Write the name of the anion, changing the “-ine” to “-ide” Watch out for polyatomic ions (just write them) Ca(OH)2 Calcium Hydroxide Calcium Hydroxide

Formula to Name Write the name of the cation If the cation is a transition metal, determine the charge (write in romans) Write the name of the anion, changing the “-ine” to “-ide” Watch out for polyatomic ions (just write them) MgBr2 Magnesium Bromine ide Magnesium Bromide

+2 -1 Formula to Name Write the name of the cation If the cation is a transition metal, determine the charge (write in romans) Write the name of the anion, changing the “-ine” to “-ide” Watch out for polyatomic ions (just write them) Total + charge: + 2 FeCl2 Total – charge: – 2 Iron II Chlorine ide Iron(II) Chloride

+3 -1 Formula to Name Write the name of the cation If the cation is a transition metal, determine the charge (write in romans) Write the name of the anion, changing the “-ine” to “-ide” Watch out for polyatomic ions (just write them) FeCl3 Total + charge: + 3 Total – charge: – 3 Iron III Chlorine ide Iron(III) Chloride

+2 -1 Formula to Name Write the name of the cation If the cation is a transition metal, determine the charge (write in romans) Write the name of the anion, changing the “-ine” to “-ide” Watch out for polyatomic ions (just write them) Total + charge: + 2 Zn(OH)2 Total – charge: – 2 Zinc II Hydroxide Zinc(II) Hydroxide

+2 -3 Total + charge: + 6 Total – charge: – 6 Formula to Name Write the name of the cation If the cation is a transition metal, determine the charge (write in romans) Write the name of the anion, changing the “-ine” to “-ide” Watch out for polyatomic ions (just write them) Co3N2 Cobalt II Nitrogen ide Cobalt (II) Nitride

Complete CW 11, CW 12, and CW 13 Complete Exit Ticket 1 HW 5: Ionic Compounds (handed out last class – check that you have a copy) Summary 8 11/8 (A Day) 11/11 (B Day) Outcome: I can write names and chemical formulas for ionic compounds Goal: CW 13

Predict the ionic formulas and write the name of the compound.
Calcium and chlorine Iron(III) and oxygen Magnesium and nitrogen Drill 9 11/12 (A Day) 11/13 (B Day) Outcome: I can explain the properties of metals using the electron sea model. Goal: CW 14, CW 15

CW 14: Naming Acids and Bases
Formula starts with hydrogen Name ends with acid Ensure that charges are neutral! Naming Common Acids Anion Ending Example Stem Acid Name -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

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 Cl-1 chloride chlor hydro____ic acid ClO-1 hypochlorite hypochlor _______ous acid ClO2-1 chlorite chlor _____ous acid ClO3-1 chlorate chlor ____ic acid ClO4-1 perchlorate perchlor _______ic acid

Name each of the following acids. H2S HCl H2SO4 HClO4 H2SO3 HClO2 sulfide Hydro___ic Acid sulf chloride Hydro____ic Acid chlor sulfate ___ic Acid Sulf perchlorate ______ic Acid Perchlor sulfite ___ous Acid Sulf chlorite ____ous Acid Chlor

Complete the table by filling in either the name or formula,. Name Formula HNO3 HNO2 HBr Hydrofluoric acid Carbonic Acid Nitrous Acid Nitric Acid nitrate Nitrous Acid nitrite Hydrobromic Acid bromide fluoride -ide HF carbonate H2CO3 -ate nitrite -ite HNO2

Name or write the formula for the following bases. NaOH Potassium hydroxide Fe(OH)2 Iron (III) hydroxide Sodium Hydroxide KOH Iron (II) Hydroxide Fe(OH)3

CW 15: Metallic Bonding Go to LEFFELlabs/ unit 2 to complete the PowerPoint while answering the questions.

CW 15: Metallic Bonding How do metals form cations? Explain using an electron configuration. Metals tend to have empty valence shells, so they get to an octet by losing e– (forms cations). Na: 1s2 2s2 2p6 3s1  Lose the 3s1 e–, 2nd energy level full Where do the electrons in the electron sea come from? Each metal atom loses e–; these form the e– sea

CW 15: Metallic Bonding Metal cation
Draw a picture that shows the electron sea model. Label: metal cation, delocalized electron, electron sea. Write a caption for your picture. Words to use: repulsion, delocalized electron, cation. Metal cation +3 Metal cations are held together by the e– sea, which prevents cations from repelling each other. Delocalized e– e– sea

CW 15: Metallic Bonding Explain each of the following in terms of the electron sea model. Metals are shiny. The e– in the e– sea easily absorb and reemit light (Bohr model, flame test) Metals are good conductors. Free flowing e– sea can carry heat and electricity Metals are malleable and ductile. The e– sea insulates metal cations from touching each other; allows cations to slide past each other

CW 15: Metallic Bonding What do the large beads represent? What do the same silver beads represent? Large beads: metal cations Silver beads: free flowing e– sea Describe the motion of the electrons relative to the cations. The electrons flow past the cations, which are immobile Use your model to define alloy. What did the replacement beads represent? An alloy is made of a mixture of metals. The replacement beads represent atoms of a different metal How might making your metal into an alloy affect the properties of the metal? Changing the structure changes properties: may become stronger/ weaker, more/ less resistant to corrosion, better/ worse conductor

CW 15: Metallic Bonding Journal Write 7
Evaluate your model. What does it illustrate well? What changes could you make to create a more accurate model?

Complete CW 14 and CW 15 HW 6: Review for Unit Test Unit test on 11/20 (B Day) Summary 9 11/12 (A Day) 11/13 (B Day) Outcome: I can explain the properties of metals using the electron sea model. Goal: CW 14, CW 15

Agenda 11/14 (A) and 11/15 (B) Review

Agenda 11/18 (A) and 11/19 (B) Unit Test

Unit 2 Journal 1 Create a table to compare the type of bonds in metalloids to those in metals (delocalized) and nonmetals (localized). Metals (Delocalized) Metalloids Nonmetals (Localized) Electrons flow freely in the e– sea Bonding is in between localized and delocalized Electrons are shared between two atoms, cannot move Shiny (e– can absorb/ emit light) May be shiny or dull Dull (localized electrons cannot absorb/ emit light) Malleable (e– allow cations to slide past each other and not repel) May be brittle or malleable Brittle (when atoms are forced together, they repel and the solid shatters) Good conductors (e– can move/ carry heat or energy with them) Semi conductors Poor conductors (e– cannot move)

Unit 2 Journal 2 Complete the summary below.
How does increasing shielding affect the attraction between the nucleus and electrons? More shielding = more electrons blocking the nucleus = weaker attraction How does increasing nuclear charge affect the attraction between the nucleus and electrons? More positive nucleus = electrons are more strongly attracted to the nucleus IN DE IN IN IN DE DE DE

Unit 2 Journal 3 and 5 Considering your answer to Question 8, why does sodium have such a low electronegativity value? Sodium has only one valence electron, which it wants to lose to get a full octet. Explain your observations of the volatility of the substances in terms of the intermolecular forces present in each molecule. The stronger the IMFs, the harder it is to convert a liquid into a gas. Stronger IMFs = less volatile. Na: 1s2 2s2 2sp6 3s1

2: Bonds, James Bonds How can we apply the physical and chemical properties of elements to predict the formation of compounds?

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2: Bonds, James Bonds How can we apply the physical and chemical properties of elements to predict the formation of compounds?

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