Chemistry Course Book Spring 2020 UNITS 7-12

7) Stoichiometry

8) Quantum Mechanics/Periodic Table

9) Bonding

10) Molecular Shapes/ Dipoles/ Intermolecular Forces

11) Thermochemistry/Thermodynamics

12) pH mini unit

Written, edited and compiled by Brian Cox

Dan Albritton

With thanks to Daniel Knowles Semester 2 Table of Contents Unit # Topic Pages Unit 7 Study Guide & 1 Assignment List 7 Notes &Problem Sets 2-13 Unit 8 Study Guide 14 8 Notes 15-22 8 Assignment List 23 8 Problem Sets 23-42 Unit 9 Study Guide 43 9 Notes 44-56 9 Assignment List 57 9 Problem Sets 58 - 74 Unit 10 Study Guide 75 10 Notes 76 - 81 10 Assignment List 82 10 Problem Sets 83-91 Unit 11 Study Guide 92 11 Notes 92-96 Problem Sets 97-118 12 pH mini unit 119-122 Semester Exam Review 123-124 Semester Exam Extra 125 Credit

Unit 7 ~ Stoichiometry ~ Test Topics 1) Meaning of Coefficients of Balanced Chemical Equations and Calculations using Balanced Chemical Equations Understand that coefficients of the balanced chemical equation represent particle ratios, NOT mass ratios. Recall that particles may be atoms, molecules, or formula units. Be able to demonstrate understanding by: • Given a balanced chemical equation, be able to draw pictorial representations of ratios of the number of reactants and products in a chemical reaction. • Given a balanced chemical equation, be able to perform calculations involving mole-mole, mass- mole, or mass-mass.

2) Limiting Reactants • Understand concept of limiting reactant (also called limiting reagent). • Be able to define limiting reactant • Be able to draw pictorial diagrams representing the concept. • Be able to determine which reactant is limiting and how much product will form.

3) % Yield • Understand that, although we can use calculations to predict the amount of product theoretically produced in a reaction, when the reaction is actually carried out in the lab less product is typically obtained than expected. • % yield is a measure of the efficiency of a reaction carried out in the lab, in other words, how much product is actually obtained compared to how much was predicted. • Know the equation for percent yield:

% Yield = Actual Yield x 100% Theoretical Yield • Given any two variables out of % yield, actual yield and theoretical yield be able to manipulate equation to solve for the third.

Unit 7 Stoichiometry Assignments

7-1 Foundation Concepts of Stoichiometry Sec. 9.1, 9.2 (p. 239 – 243) 7-2 Mass – Mass. Sec 9-3 (p. 243 – 251) 7-3 % Yield Section 9.5 (p. 257- 258) 7-4 Limiting Reactant Concept 7-5 Limiting Reactant Calculations Section 9.4 (p. 251 – 257) 7-6 Conceptual Review 7-7 Unit 7 Review Problem Set

Foundation Concepts of Stoichiometry: HW 7-1

Unit 7 Notes ~ Stoichiometry Stoichiometry – is the study of the relationship between the ______of ______and ______in a chemical reaction. Importance: One the most important task of chemists is to make compounds with desirable properties. For example, antibiotic drugs, anti-cancer medications, plastics, computer chips, etc. Stoichiometry provides ______for chemical reactions. Key Concept of this unit:

Part 1: Interpreting the Coefficients of a Balanced Chemical Equation For questions #1-3, use symbols to represent the relative number of particles in each of the following equations:

1) 2 Mg + O2 → 2 MgO Use Mg = O =

2) 2 H2 + O2 → 2 H2O Use H = O =

3) N2 + 3 H2 → 2 NH3 Use N = H =

For questions, #4-6, use the balanced chemical equation to fill in the missing number of particles (or moles) in the blanks provided:

4) Balanced Equation: 2 Mg + O2 → 2 MgO a) Given: 4 atoms + 2 molecules → ____ molecules

b) Given: 8 atoms + ___ molecules → 8 molecules

c) Given: ___ atoms + 10 molecules → _____ molecules

5) Balanced Equation: N2 + 3 H2 → 2 NH3 a) Given: ______molecules + 15 molecules → ____ molecules b) Given: 6.022 x 1023molecules + ______molecules → ______molecules

C) Given: 1 mole + ______moles → ______moles

6) Balanced Equation: 2 H2 + O2 → 2 H2O

A) Given: 2 moles + ______moles → ______moles

B) Given: ____ mole + 4 moles → ______moles

C) Given: ____ mole + ____ moles → 10 moles

Part 2: Complex Mole Ratios Using Factor Label Dimensional Analysis Use dimensional and show your work.

7) Balanced Equation: 4 NH3 + 7 O2 → 4 NO2 + 6 H2O

A) How many moles of NH3 would be needed to completely with 10.0 moles of O2?

10.0 moles O2 =

( )

B) How many moles of H2O would be produced when 11.5 moles of NH3 are reacted?

11.5 moles NH3 =

( )

C) How many moles of NO2 would be produced when 8.6 moles of H2O are produced?

D) How many moles of H2O would be produced when 9.2 moles of O2 are reacted?

Mass-Mass Homework 7-2 Notes~ Mass-Mass Problems Problem: COEFFICIENTS of the balanced chemical equation represent MOLE RATIOS but in the lab the only instrument we have for measuring amounts of atoms are scales. Solution: Mass-Mass calculations require you to link the mole concept idea of counting by weighing with the mole-ratio concept.

Example: 4 Fe + 3 O2 → 2 Fe2O3 167.4 g ? g

How many grams of Fe2O3 could be produced from the reaction of 167.4 g of Fe?

167.4 g Fe = ( )( )( )

You must show work using factor label to receive credit.

1) 4 Al + 3 O2 → 2 Al2O3

Molar Masses: Al: 26.98 g/mole O2: 32.0 g/mole

You wish to prepare 153 g of Al2O3: Al2O3 = 101.96 g/mol

A) Calculate the number of grams of Al needed to prepare 153 g of Al2O3

153 g Al2O3 =

( )( )( )

B) Calculate the number of grams of O2 needed to prepare 153 g of Al2O3

153 g Al2O3 = ( )( )( )

2) N2H4 + 3 O2 → 2 NO2 + 2 H2O

Calculate Molar Masses: N: N2 : H2 : To use in parts A + B O2: H4 : O:

Total Molar Mass: NO2: N2H4: H2O:

A) How many grams of NO2 could be produced from the reaction of 64.0 g of N2H4?

B) How many grams of H2O could be produced from the reaction of 64.0 g of N2H4?

3) C2H6O + 3 O2 → 2 CO2 + 3 H2O

A) How many grams of CO2 will be released by the combustion of 50.0 g of C2H6O?

B) How many grams of H2O will be released by the combustion of 92.0 g of C2H6O?

4) CO + 2 H2 → CH3OH What mass of carbon monoxide is required to produce 6.0 kg of CH3OH?

Answers:

1A) 81.0 g Al 1B) 72.0 g O2 2A) 184 g NO2 2B) 72.0 g H2O 3A) 95.6 g CO2 3B) 108 g H2O 4) 5.3 x 103 g CO

Percent Yield Notes:

% Yield = ______x 100%

Theoretical yield- calculated (grams → moles → MOLES → GRAMS) amount of product Theoretical yield - Actual yield -

Example: K2CO3 (s) + 2HCl(aq) → 2 KCl(s) + H2O(l) + CO2(g)

Calculate the theoretical yield of KCl solid if 45.8 g of K2CO3(s) are reacted.

The actual amount of KCl recovered in the lab is 46.3 g, less than the theoretical yield. % Yield =

Importance of % of Yield:

HW 7-3 1) Student’s Prelab theoretical yield calculation = 1.44 g MgO. When student weighed product in lab after reaction, only 1.23 g MgO obtained. Calculate % yield.

2) 2 HgO → 2 Hg + O2 a) When 1.25 g of HgO is heated, what is theoretical yield of mercury, Hg?

b) Suppose 1.09 g of Hg is actually collected. What is % yield?

Answers: 1) 85.4 % 2) theoretical yield: 1.16 g; % yield: 94.0% 6

Concept of Limiting Reactant HW 7-4

Limiting Reactant Conceptual Notes

Mg + I2 → MgI2 Draw an atomic level picture of the following reaction using Mg = and I = Identify the limiting reactant and the reactant which is present in excess. Write None if neither reactant is limiting.

A) 1 atom Mg + 1 molecule of I2 → ______molecules MgI2

Limiting: Excess:

B) 1 atom Mg + 2 molecule of I2 → ______molecules MgI2

Limiting: Excess:

C) 3 atom Mg + 2 molecule of I2 → ______molecules MgI2

Limiting: Excess:

Conceptual Practice: Draw an atomic level picture of the following reaction. Identify the limiting reactant and the reactant which is present in excess. Write None if neither reactant is limiting.

1) 2 H2 + O2 → 2 H2O Use H = O =

A) 2 molecules H2 + 2 molecules of O2 → ______molecules of H2O

Limiting Reactant ______Excess Reactant ______

B) 4 molecules H2 + 2 molecules of O2 → ______molecules of H2O

Limiting Reactant ______Excess Reactant ______

C) 3 molecules H2 + 1 molecules of O2 → ______molecules of H2O

Limiting Reactant ______Excess Reactant ______

Notes ~ Limiting Reactant Calculations Key Concept: In these problems you will have two gram amounts (one for each reactant). You must decide, by doing calculations, which gram amount will “run out first” and stop the reaction!

Problem: A fuel mixture used in the early days of rocketry is composed of two liquid, hydrazine (N2H4) and dintrogen tetroxide (N2O4), which ignite on contact to form nitrogen gas and water vapor. These hot, expanding gases push the rocket up.

2 N2H4 + N2O4 → 3 N2 + 4 H2O 100. g 200. g ? g

How many grams of N2 gas are formed when 100. g of N2H4 and 200. g of N2O4 are mixed?

Step 1: Calculation of grams N2 produced from reaction of 100. g of N2H4:

Step 2: Calculation of grams N2 produced from reaction of 200. g of N2O4:

Final Answer: Explanation:

HW 7-5 – Limiting Reactant Calculations (from Zumdahl self-check 9.6, 9.7, p. 265 #65 and p.266, # 69) 1) 6 Li + N2 → 2 Li3N Calculate mass of lithium nitride formed from 56.0 g nitrogen gas and 56.0 g of lithium. 56.0 g 56.0 g ? g (Self-check 9-6) Calculation of g Li3N formed from 56.0 g of Li:

Calculation of g Li3N formed from 56.0 g of N2:

Final Answer: Explanation:

2) TiCl4 + O2 → TiO2 + 2 Cl2 ( Self-check 9.7) 3 3 a) 6.71 x 10 g TiCl4 reacts with 2.45 x 10 g oxygen. Calculate maximum mass of TiO2 that can form.

b) If the % yield for the reaction = 75%, what mass of TiO2 actually formed?

3) Xe + 2 F2 → XeF4 (p. 265, #65) a) 130 g 100. g ? g theoretical yield

b) Actual yield = 145 g XeF4 , calculate % yield.

4) Fe + S → FeS (p. 266, #69 ) 5.25 g + 12.7 g ? g theoretical yield

3 3 Ans:1) 9-6: 93.7 g Li3N; 2) 9-7: TY =2.83 x 10 TiO2 ; AY = 2.12 x 10 TiO2 3) TY = 205 g; % yield = 70.7 % 4) Fe limiting: 8.26 g FeS

HW 7-6 Conceptual Review

C3H8 + 5 O2 → 3 CO2 + 4 H2O Use your knowledge of ratios to fill in the missing information in the table below: Note: XS = reactant present in excess; thus other reactant = limiting

C3H8 O2 CO2 H2O 1 molecule 5 molecules molecules molecules 1 molecule XS molecules molecules 2 molecules 5 molecules molecules molecules 3 molecules XS molecules molecules 6.022 x 1023 molecules (6.022 x 1023 molecules) (6.022 x 1023 molecules) 4 (6.022 x 1023 molecules) 1 mole moles moles 4 moles moles 10 moles moles 8 moles XS moles 6 moles 8 moles 0.5 moles 5 moles moles moles moles 1.0 moles 0.6 moles moles 4 moles XS moles moles 2 moles 11 moles moles moles

Molar Masses: C3H8 = 44 g/mole O2 = 32 g/mole CO2 = 44 g/mole H2O = 18 g/mole 44 g 5 (32 g) = 160 g 3 (44 g) = 132 g (18 g) = 44 g 200 g (44) = g (18 g) = g 22 g 160 g (44) = g (18 g) = g

2) Given the following reaction: 2 P + 3 Cl2 + → 2 PCl3

A) Using to represent P atoms and to represent Cl atoms, draw a diagram to represent the balanced chemical equation above.

B) Draw a diagram to represent the reaction of 3 atoms of P with 3 molecules of Cl2. Which reactant is limiting? Which reactant is in excess and how much would be left unreacted?

C) Draw a diagram to represent the reaction of 2 atoms of P with 4 molecules of Cl2. Which reactant is limiting? Which reactant is in excess and how much would be left unreacted?

D) How many molecules of Cl2 would be needed to complete react with 4 atoms of P?

HW 7-7: Stoichiometry Review Problem Set

1) CH4 + 2 O2→ CO2 + 2 H2O (Methane mamba demo!)

For each combination of CH4 and O2,, indicate how many product CO2 molecules, H2O molecules and leftovers would be produced.

# CH4 molecules O2 molecules # CO2 molecules # H2O molecules Reactant Leftovers 1 2 2 2 4 9 5 11 8 13

2) Concept check of limitations of comparing small # of molecules vs moles: Given the following reaction:

A + B → C A) The stoichiometry of the reaction is 1:1. Is it physically possible to half the recipe by reacting ½ of an atom of A?

B) Is it physically possible to half the recipe by reacting ½ mole of atom of A?

MUST SHOW WORK ON CALCULATION PROBLEMS TO EARN CREDIT!

3) CH4 + 2 O2→ CO2 + 2 H2O (Methane mamba demo!) A) To produce 1.0 mole of carbon dioxide and 2.0 moles of water, how moles of each reactant are needed?

B) If 8.0 moles of methane are available for combustion, how many moles of water will be produced?

C) If 8.0 moles of methane are available for combustions how many grams of water will be produced?

D) If 128.4 g of methane are available for combustion, how many grams of water will be produced?

4) Given the following reaction: 2 C2H6 + 7 O2 → 4 CO2 + 6 H2O

A) How many moles of CO2 could be produced when 14 moles of O2 react with excess C2H6?

B) How many moles of H2O could be produced when 3.5 moles of O2 reacts with excess C2H6?

C) If 4 moles of C2H6 are mixed with 19 moles of O2, which reactant would be in excess and how many moles would be unreacted?

D) If 3 moles of C2H6 are mixed with 14 moles of O2, which reactant would be in excess and how many moles would be unreacted?

5) Given the following reaction: 2 Al + 6 HCl → 3 H2 + 2 AlCl3

A) If 108.0 g of Al reacts with excess HCl, how many moles of H2 could be formed?

B) ) If 108.0 g of Al reacts with excess HCl, how many grams of H2 could be formed?

6) Chlorine gas is a very reactive substance and will combine with most metals. One reaction is with aluminum:

2 Al + 3 Cl2 → 2 AlCl3

A) How many moles of product, AlCl3, are formed from 1.32 moles of Cl2 gas?

B) If 25.0 g of Al are mixed with 50.0 g of Cl2, determine whether the metal or the chlorine is the limiting reactant and calculate the theoretical yield of the metal chloride.

7) Ammonia, NH3 reacts with oxygen, O2 in the following reaction:

4 NH3 + 5 O2 → 4 NO + 6 H2O

How many grams of NO could be formed by the reaction of 2.00 g of NH3 with 4.00 g of O2? What is the limiting reactant in this problem?

8) When the sugar glucose, C6H12O6, is burned in air, carbon dioxide and water vapor are produced. The balanced chemical equation for this process is:

C6H12O6 + 6 O2 → 6 CO2 + 6 H2O A) What is the theoretical yield of carbon dioxide when 1.00 g of glucose is burned completely?

B) What is the percent yield if 1.25 grams of CO2 are actually obtained when carrying out the reaction in the lab?

9) Before going into the lab, a student read on his lab sheet that the percent yield for a difficult reaction to be studied was likely to be only 40.0 %. The student’s prelab stoichiometric calculations predict that the theoretical yield should be 12.5 g. What is the student’s actual yield likely to be?

Answers to selected problems: 3A) 2 mol O2 , 1 mol CH4 3B) 16 mol H2O 3C) 290 g H2O 3D) 288.5 g H2O 4A) 8 moles CO2 4B) 3 moles H2O 4C) 5 moles O2 excess 4D) 3.5 moles O2 excess 5A) 6.004 moles H2 5B) 12.13 g H2 6A) 0.880 moles AlCl3 6B) Cl is limiting; 62.7 g AlCl3 formed 7) O2 is limiting; 3.00 g of NO are formed. 8A) 1.46 g CO2 8B) 85.6 % 9) 5.00 g

Unit 8 ~ Quantum Mechanics/the Periodic Table ~ Test Topics

Electromagnetic Radiation

• Understand terms wavelength, frequency and amplitude of electromagnetic radiation • Understand that wavelength and frequency determine the type of electromagnetic radiation. High frequency (short wavelength) = high energy radiation. • Different colors of visible light have different wavelengths and therefore different energies. Relative energy of visible light per photon increases from red → orange → yellow → green→ blue→ violet Atomic Models

• Be able to discuss key ideas, strengths, weaknesses, and the supporting experimental evidence for the following atomic models: Rutherford, Bohr, Modern Quantum Mechanical (current model) Quantum Concept • Understand the concept of quantized energy levels and be able to explain how the observed emission spectrum of hydrogen supports the idea of quantized energy levels for electrons. Foundations of Modern Quantum Theory • Understand implications of Heisenberg Uncertain Principle for describing electron motion: we cannot trace the exact path of the electron; we can only describe the probability of finding an electron in a particular region of space (orbital). Orbitals • Know definition of orbital: 3-D region of space in which there is a 90% chance of finding an electron. • Know the difference between a Bohr “orbit” and the “orbital of modern quantum theory. • Know that average distance the electrons are from the nucleus increases with a higher energy level. (Example: what is the difference between 1s, 2s and 3s orbitals.) • Be able to sketch general shapes for s, p and d (4 leaf clover) orbitals. Electron Configurations • Be able to use the periodic table to determine the electron configuration for elements in the first 5 periods of the periodic table. • Be able to write both full electron configurations and orbital notation diagrams (box/circles and arrows or core abbreviation notation with orbital notation for valence electrons only. • Understand the relationship between position on periodic table, electron configuration, and chemical behavior. Valence electron configuration determines chemical behavior.

Periodic Trends • Know trends in atomic radius, ionic radius and ionization energy on the periodic table and be able to link to chemical properties of elements. • Successive Ionization Energies – be able to analyze successive ionization energy data (looking for jump from “easily” removable valence electrons to “difficult to remove” filled valence electrons) to determine the number of valence electrons.

Introduction to Quantum Mechanics ~ Notes

• Study of the behavior ( ______and ______) of electrons within the atom. • Importance: Energy arrangement of electrons within atom determines ______(how atoms react chemically). Sketch Rutherford Model Atom in box below

Rutherford Model of the Atom • Atom contains a ______surrounded by an electron cloud. • Nucleus is composed of ______and ______. • Motion of electrons is ______. Questions Raised by Rutherford Model:

• If no restrictions are placed on the motion and energy of the electron

Sketch path of e- if no restrictions The Atom is Stable! on motion and energy Since the electron is into the nucleus is was very clear to scientists in 1911 that there were very important properties of the electron that no one understood! Question: How do you investigate the behavior of electrons inside the atom?

Answer: By investigating the interaction of

Nature of Light

• Light is a form of ______.

• All forms of electromagnetic radiation travel at the speed of light 3.00 x 108 m/s or 186,000 mi/s.

• Light can be visualized as particles, little packets of energy called ______.

• Light can also be thought of as ______. Some of the properties of light can be understood in terms of wave behavior.

Wavelength ( ) and Amplitude • Wavelength = distance between 2 wave ______or ______.

• Amplitude = ______of wave.

Frequency

• Frequency = ______that pass by a fixed point per second.

• As wavelength ______, frequency ______.

• Frequency is measured in ______(s-1).

Classification of Electromagnetic radiation by Energy

• The type of electromagnetic radiation is determined by its energy. • Energy of an electromagnetic wave depends ______and ______.

Electromagnetic Spectrum of Visible Light (Highest to Lowest Energy)

Violet Red

Wave Interference

• When waves interact (collide), the amplitudes can either add or subtract. This is called interference.

• Constructive Interference: wave amplitudes ______.

• Destructive interference: wave amplitudes ______.

Constructive Interference Destructive Interference Before Interference During Interference Before Interference During Interference

Double Slit Experiment

• Light projected on a screen through 2 slits shows dark areas where light disappears and bright areas where intensity of light increases. • Result can only be explained by concept of wave interference.

Quantum Mechanical Interpretation of Emission Spectra of H2 Experimental Observation: ADD ______(in the form of electricity, heat or light) Specific sharp lines of color are produced. Each color of light has a specific wavelength hence a specific energy.

KEY OBSERVATION: ONLY CERTAIN, ______OF LIGHT ARE BEING RELEASED.

Energy release is ______= only certain specific energies are being emitted.

Interpretation of the Experiment: Energy added to atoms is absorbed by the ______. When an electron absorbs a specific quantity of energy, it makes a quantum leap from its lowest energy state (______) to a higher energy state (______).

n=3 ______excited state n=3 ______excited state

n=2 ______excited state n=2 ______excited state

n=1 ______ground state n=1 ______ground state

n = 1 to n =2 quantum leap n = 1 to n =3 quantum leap

Heat, Light, Electricity (Energy Added) → Electron makes quantum leap from ground state to excited state

= electron; Use arrows to show quantum leaps.

Note: Only certain distance leaps are possible, for example, n=1 to n=2 or n=1 to n=3 or n=2 to n=3.

Electrons ______remain in an excited energy level for any significant period of time. The electron in the excited state returns to the ground state by emitting a specific amount of energy, in the form of ______. The energy of light emitted represents the exact difference in energy between the excited state and the ground state.

Key Point:

n=3 ______excited state n=3 ______excited state n=3 ______excited state

n=2 ______excited state n=2 ______excited state n=2 ______excited state

n=1 ______ground state n=1 ______ground state n=1 ______ground state

n = 3 to n =1 qua ntum leap n = 3 to n =2 quantum leap n = 2 to n =1 quantum leap

When an electron, , makes a quantum leap down from an excited state to a lower energy state, ENERGY DIFFERENCE RELEASED AS LIGHT. Different energy differences result in the emission of different colors of light.

Note: Only certain distance leaps are possible, for example, n=3 to n=1 or n=3 to n=2 or n=2 to n=1;

CANNOT have a jump from 3 to 1 ½ for example. Note also that the energy difference between 2 to 1 is different than the energy difference from 3 to 2 so that each transition gives off a different energy of light.

Foundations of Modern Quantum Mechanics Bohr model (1913) • Concept of ______(restricted, only certain specific allowed) energy levels.

• Concept of electron level transitions or ______accurately predicts visible emission spectra of hydrogen, i.e red line at 656 nm, teal line at 486 nm, blue line at 434 and violet line at 410 nm.

• Bohr model does NOT accurately predict emission of spectra of any element besides hydrogen. • Bohr Model FATAL FLAW: Electrons DO NOT travel in fixed ______.

• Modern Quantum Mechanics ______the concept of quantized energy levels, but ______Bohr’s circular orbits.

Question: Is it possible to determine the exact trajectory of an electron around the nucleus? Answer:

Heisenberg Uncertainty Principle (1920’s) It is impossible to accurately determine both the position and momentum (mass x velocity; note velocity = speed and direction) of an electron at the same time.

Explanation of Uncertainty Principle: • “Seeing” electron requires use of very short wavelength, high energy photon of light • The photon collides with the electron; the scattered photon reports the position of the collision, however the energy transferred to the electron as the result of the high energy collision causes the electron to move off in a new direction/new speed which cannot be measured or calculated. • The act of measuring the ______changes the ______of the electron.

Implication of Uncertainty Principle: It is ______to trace the exact trajectory (path) of an electron.

It is only possible to describe the ______of the electron being in a particular place at a particular time.

Developing Probability Maps of Electron’s Location Using Wave Functions DeBroglie Wave-Particle Duality (1920’s) • All matter has properties of both particles and waves. • Combined famous equations describing energy in terms of mass (a particle property), E = mc2 and frequency a property of waves, E = hv, to predict wavelength of electron. • The wave properties of matter become very important when considering very small mass particles (such as an electron or a photon of light) moving at very high velocities (at or near the speed of light).

Schroedinger Equation • Developed very complex equations that described the electron as standing wave around nucleus. • Recall that waves can interfere constructively or destructively. Destructive interference results in collapse of wave function. • Results explain why electron energies are ______. • Only certain specific ______are possible for a standing wave to “fit” in the space around the nucleus without destructive interference collapsing the wave function.

• Solutions to wave function equations allow calculations of the ______of finding an electron in a particular region of a space. • Probability map of an electron within an atom is called an orbital.

ELECTRON ORBITAL NOTES

Modern quantum theory ______(orbit) of an electron. The modern wave mechanical model describes a general area of ______where there is a ______of finding an electron.

Orbital Orbit

Defn of Orbital: A region of ______(3D) in which there is

Capacity of Orbital: Each orbital can hold a maximum of _____ electrons (1 pair).

Orbital notation diagrams: An individual orbital can be represented as or or _____

______Use: to represent electrons.

4 General Types (or shapes) of Orbitals: (On test be able to draw general shape of s, p, d; Know number of sets each type of orbital (s,p,d,f) per energy level.

s orbital p orbital d-orbital

Shape: Shape: Shape: f orbital

Number: Number: Number: Number:

Energy levels: 1-7 Energy levels: 2- 7 Energy levels: 3-7 Energy levels: 4-7

Total Electron Capacity of Each Sublevel: (Number in orbital set x 2 e- / orbital ) s: 1 x 2 = 2 p: x 2 = d: x 2 = f: x 2 =

Relationship between Energy Level and Orbital Size: In general: The ______the energy level, the ______the ______the electron is from the nucleus.

Higher Energy Level = (on average e- are farther from nucleus) 1s 2s 3s

Electron Configurations and Orbital Notation Diagrams Electron Configuration – describes the orbitals electrons occupied within the Atom; helps predict chemical behavior Electron Configuration: Hydrogen 1s1

Orbital Notation: Represent orbitals as or or And electrons as

___ orbital; ____ per energy level:

____ orbitals; ____ per energy level:

Experimentally observed order of filling orbitals: Recall each orbital can hold up to ____ electrons.

Sublevel electron capacities:

Type Levels present Number x 2e- Max Capacity

Aufbau rules • Fill ______energy orbitals first.

• Each orbital can hold ______electrons with ______spins.

Hund’s Rule • Within a sublevel, place _____ electron per orbital spin up before pairing (Placing 2nd e- spin down).

Unit 8 Assignments – Quantum Mechanics

8-1 Sec. 11.1, 11.2 (p. 303 – 307) Introduction to Quantum Mechanics

8-2 Sec. 11.3, 11.4 (p.308 – 311) Hydrogen Emission and Bohr Model

8-3 Section 11.5 (p.311 – 312) Foundations of Modern Quantum Mechanics

8-4 Sec. 11.6-11.8 (p. 312 – 318) Electron Orbitals

8-5 Quiz Review Unit Review Part 1/ Practice Quiz

8-6 Sec. 11.9, 11.10 (p. 319-327) Electron Configurations and Orbital Notation

8-7 Core Abbreviation

8-8 Electron Configurations and Chemical Reactivity

8-9 Successive Ionization Energies

8-10 Atomic Size

8-11 Atomic Size Trends

8-12 Alkali Metals – Chemistry of a Family

8-13 Unit 8 Review Problem Part 2

Note: Homeworks 8-1 thru 8-5 will be turned in at the time of our Quantum Quiz. (Covers in Zumdahl sections 11.1 thru 11.8). Homeworks 8-6 thru 8-16 will be due at the time of the Unit Test HW 8-1- Introduction to Quantum Mechanics 1) Rutherford’s experiment demonstrated that the atom has a small _____ charged core called the nucleus with ______charged electrons moving in space around the nucleus.

2) What questions were left unanswered by Rutherford’s experiments?

Review of Electromagnetic Waves 3) What is electromagnetic radiation? At what speed does electromagnetic travel?

4) Does light consist of waves, or is it a stream of particles of energy? Or is it both? Explain.

5) What do we mean by the frequency of electromagnetic radiation? Is the frequency the same as the speed of the electromagnetic radiation?

6) How are the different types of electromagnetic radiation similar? How are they similar?

7) Rank the following colors of light from highest to lowest energy of light per photon: Blue, Green, Red, Purple

Review of Waves: Wave A Wave B

8) Which wave has a larger wavelength ( ) ? _____

9) Which wave has a higher frequency? ______

10) Which wave has higher energy (E), (hint: which has more “ups and downs”) ______

11) From your answers above, which variable is directly proportional to energy (as one goes up, the other goes up too), wavelength or frequency? ______

12) If the time pictured is one second, what is the frequency of wave A in Hertz? ______Hz

13) If the time pictured is one second, what is the frequency of wave B in Hertz ______Hz

High Frequency, short wavelength High Frequency, short wavelength

14) Which has higher energy, ultraviolet light or visible light?

HW 8-2 –Bohr Model and Concept of quantized energy levels in the Hydrogen Atom

1) What are the essential points of Bohr’s theory of the structure of the hydrogen atom?

2) What does it mean to say that the hydrogen atom has quantized energy levels?

3) What does the ground state of an atom represent?

4) When an electron in an excited state returns to its ground state, what happens to the excess energy?

5) What determines the energy of light emitted when an electron makes a quantum leap from excited state to ground state?

6) When high voltage is applied to a tube of hydrogen gas, why are only certain specific colors of light emitted instead of all colors of light?

HW 8-3 – Foundations of Modern Quantum Theory

1) Why was the Bohr model initially accepted (i.e. what experimental observations did the model explain), and then why was it ultimately rejected (i.e. what experimental observations was the model unable to explain)?

2) According to the Heisenberg Uncertainly Principle, if we try to locate the exact position of an electron by hitting it with a high energy photon, what will happen to the velocity (speed and direction) of the electron?

3) Is it possible to ever determine, either by a calculation or an experimental measurement the exact path of an electron within the atom? How does modern quantum wave theory describe electron motion?

HW 8-4 Electron Orbitals

1) Discuss the difference between an orbit (as described by Bohr for hydrogen) and an orbital (as described by the more modern, wave mechanical picture of the atom).

2) Explain the “firefly” analogy helps us understand the modern wave mechanical model of describing the behavior of the electron.

3) What is an electron orbital?

4) What is the maximum number of electrons that can occupy any single orbital?

5) Sketch the following: i) s orbital ii) p orbital iii) d orbital

6) Fill in the missing information in the table below:

Orbital type Energy levels # of orbitals per energy Total number of electrons level (Comes as a set of ) sublevel can hold:

d 3 thru 7

f 4 thru 7

7) On the periodic table below label the s, p and d blocks: 8) Sketch a 1s orbital next to a 2s orbital. Explain how 1s and 2s orbital are similar and how they are different.

1s 2s .

HW 8-5 Unit Review Part 1 Practice Quantum Quiz

Note: Each wave has the same height

1) Which wave i,ii or iii has the highest frequency?

2) Which wave i,ii or iii has the longest wavelength?

3) Which wave i,ii or iii has the highest amplitude?

4) Which wave i,ii or iii has the highest energy?

5) If the waves are electromagnetic waves, which wave i,ii or iii is traveling at greatest velocity?

6) Draw a labeled diagram of the Rutherford Model of the atom.

7) According to the laws of physics, if an electron has no restrictions on its motion and energy, why would the atom be unstable? (What would happen to the electron?)

8) What important assumption did Bohr make about the motion and energy of the electron within the atom?

9) What is the word we use to describe the energy and motion of an electron which means restricted values; only certain specific (discrete) values are possible?

10) When high voltage electricity is passed through a glass tube filled with hydrogen gas, 3 colors of light red, teal and purple are emitted. A) Explain, using both words and a diagram what is happening inside the atom to produce the light emissions.

B) The electronic transitions that produce the light are 5→ 2, 4→ 2 and 3 → 2. Which transition corresponds to which color of light and how do you know?

11) What idea from the Bohr model do we no longer believe?

12) What is the implication of the Heisenberg Uncertain Principle for describing the motion of electrons?

13) What is an orbital?

14) What is the maximum number of electrons an orbital can hold?

15) Sketch the following orbitals: i) s ii) p iii) d

16) Sketch a 2p and 3p orbital side by side. How are they alike and how are they different?

17) For each orbital type, s, p,d and f indicate the number of orbitals per energy level (how many are in a set?) and indicate how many total electrons can be held in the sublevel.

s: # orbitals: p: # orbitals: d: # orbitals f: # orbitals

total e- sublevel:

HW: 8-6: Electron Configurations and Orbital Notation Diagrams Write the electron configurations and orbital notation diagrams for the first 8 elements: Electron configuration Orbital Notation diagram H, #1: 1s1

He, #2: 1s2

Li, #3: 1s22s1

Be, #4:

B, #5:

C, #6:

N, #7:

O, #8:

p. 321, Self-check 11.2: Write complete electron configurations and orbital notation diagrams for each of the elements Al (#13) thru Ar (#18).

Al (#13)

Si (#14)

P (#15)

S(#16)

Cl (#17)

Ar (#18)

Homework 8-7 Core Abbreviation

Core electrons - ______electrons; these levels are completely filled.

Valence electrons- ______electrons; these levels are the last levels in the process of being filled.

Importance of Valence Electron Configuration: Determines ______

Complete electron Configuration for Na#11: Na: 1s22s22p63s1

Core

Core abbreviation: 1s22s22p6 = [Ne];

Core abbreviation for Na#11 = [Ne]3s1

*************************************************************************************************

Write core abbreviations electron configuration for the following elements. Write orbital notation diagrams for valence electrons only.

1) P, #15 2) Ar, # 18

3) Ca, #20 4) V, #23

5) Fe, #26 6) Se, #34

7) Zr, #40 8) Sn, #50

HW 8-8 Electron Configurations and Chemical Reactivity

Electron configurations help us predict detailed chemical and physical properties. Chemical Properties Key Idea: Certain valence electron configurations such as filled valence electron levels are very stable.

Atoms gain, lose or share electrons to optimize their valence electron configurations.

Part 1: Noble Gas Family:

1) Write the complete electron configuration for Helium (element #2).

2) Write the core abbreviation electron configuration for Neon (element #10), Argon (element #18) and Krypton (element #36).

Ne: Ar: Kr:

3) Neon, Argon and Krypton all share common pattern in their s and p sublevels. What is the pattern?

4a) The Noble Gases are chemically inert (unreactive). What characteristic of their valence electron configuration explains why this family of elements are unreactive?

4b) Why is Helium placed in the same family as Neon, Argon and Krypton even though it does not have the ns2np6 (where n = the energy level 2 - 7) pattern as part of its electron configuration? (Hint: What sublevels are available for the first energy level, n=1?)

Predicting chemical reactivity for Alkaline Earth Family:

5A) Write a core abbreviation electron configuration for Ca (element #20) and Sr (element #38).

5B) In its chemical reactions Ca and Sr typically form +2 ions. Which two electrons were lost? Explain, in terms of valence electron configuration, why Ca+2 and Sr+2 ions are so stable.

5C) Give two examples of tissues in our bodies that contain high concentration of calcium.

5D) In 1986 a fission nuclear reactor in Chernobyl, Ukraine experienced a steam explosion that released radioactive isotopes into the atmosphere including Sr-90 a radioactive product of nuclear fission. Ingestion of Sr-90 is believed to have caused many cases of cancer among people living and working near the Chernobyl site. Sr-90 undergoes beta decay (emission of an electron from the nucleus.). Sr-90 has been shown to concentrate in bone and cause leukemia by radiating bone marrow. Propose an explanation for why Sr-90 is concentrated in bone.

Using Electron Configurations to gain insight into the Chemistry of Transition Metal Ions

Recall from Unit 2 that transition metal chemistry is much more varied and complex than, for example, Alkaline Earth metals. However, the concept of stable electron configurations can offer some insight into the charges on transition metals.

6A) Write the core abbreviation for Zinc, element #30.

6B) The core abbreviation for the stable ion of Zn, Zn+2 is [Ar]3d10. Compare this configuration to configuration for the neutral element you wrote in 6A. Which electrons have been lost?

6C) Concept of “stable electron configurations” for transition metals

The core abbreviation electron configurations for the ions Ag+1, Zn+2 and Cd+2 are

Zn+2 : [Ar]3d10

Ag+1 : [Ar]3d10

Cd+2 : [Kr]4d10

What electron configuration is present in all three and suggests one possible stable pattern for transition metal ions?

HW 8 -9 ~ Successive Ionization Energies and Valence Electrons

Ionization Energy – is the energy required to remove an electron from an atom in the gas phase. Successive ionization energies refers to the energy to remove multiple electrons from an atom, in order one at a time. For example, the first ionization energy, IE1 is the energy required the first electron from an atom, the second ionization energy, IE2, is the energy required to remove the second electron and so on. Consider the following data:

st nd rd Element IE1: 1 Ionization Energy IE2: 2 Ionization Energy IE3: 3 Ionization Energy (kJ/mol) (kJ/mol) (kJ/mol) Li 520 7,297 11,810 Be 900 1,757 14,480 Ne 2,080 3,963 6,276 Na 496 4,565 6,912 Mg 738 1,450 7,732 Ar 1,520 2,665 3,947

1) Which elements have very high 1st high ionization energies? ______and ______

• Write the valence electron configuration for each element.

• How can you explain the relative ionization energies in terms of the valence electron configurations?

2) Which elements have relatively low 1st ionization energies but relatively high 2nd ionization energies? ______and ______

• Write the valence electron configuration for each element.

• How can you explain the relative ionization energies in terms of the valence electron configurations?

3) Which elements have relatively low 2nd ionization energies but relatively high 3rd ionization energies? ______and ______

• Write the valence electron configuration for each element.

• How can you explain the relative ionization energies in terms of the valence electron configurations?

HW 8-10: Atomic Size Problem Set

1) What 3 particles make up an atom?

2) For each of the following 3 atoms, indicate the number of protons, neutrons and electrons present. 7 Li 20Ne 23Na

3 10 11

p; p; p;

n: n: n:

e-: e-: e-;

3) A reasonable first assumption about atomic size is that atoms that contain more particles (protons, neutrons and electrons) would be the largest atoms, and that particles with the fewest number of particles would be the smallest. Based on this assumption, rank the atoms in order from largest size to smallest size.

Structure of the atom – what does atomic size really mean? 4) At the beginning of the course, we discussed the Rutherford scattering experiment. Recall that Rutherford fired positively charged alpha particles at gold atoms. a) Complete the sentence: The vast majority of alpha particles

b) What does this result tell about the relative volume of the nucleus compared to the volume of the electron cloud? Sketch an atom showing the relative size of the nucleus compared to the electron cloud.

5) Fill in the blank: Since the nucleus is so tiny compared to the total volume of the atom, the size of the atom is really a measure of the size of the ______cloud.

6) Which atom Li, Ne or Na has the greatest number of electrons? The fewest? Based on the number of electrons, predict the relative size of the atoms.

Actually Observed Sizes

7) The actually observed order of sizes is: Na (186 pm) > Li (152 pm) > Ne (71 pm).

How can we explain the observation that Li is bigger than Ne and Na is so much larger than Ne even though Na and Ne have almost the same number of protons and electrons? In order to make sense of these observations, answer the series of questions below.

a) Write the electron configuration of each element: Li, #3: Ne, #10: Na, #11:

b) Even though the Bohr model is an oversimplification, the concept of the energy levels is useful for our analysis. Using the Li diagram below as an example, complete the picture for Na and Ne. (Hint: Place both the 2s and 2p electrons together in the second ring).

Li, (3 protons, 2 electrons in first level, 1 e- in 2nd level)

c) Li vs Ne: Period trend (across a row). Look at your diagrams Li has 1 electron in the 2nd (valence (level, while Ne has 8 in its second (valence) level. What other factor could cause Neon’s electrons to move closer to the nucleus than Li electrons?

d) Na vs Ne. Na and Ne only differ by 1 proton and 1 electron (11 vs 10) and yet Na is more than twice as large as Ne (186 pm vs. 71 pm). Look at the pictures you drew representing the electrons in their level. What is the key difference that explains Na much larger size?

Homework 8-11 Periodic Trends in Atomic Radius

Atomic radius – measures the size of the atom

K Rb Na Li aa a

Ar Kr He Ne a a aa a

He a Li Ne

Na Ar aa aa K Kr a Rb a

1) On the periodic table, horizontal rows are called ______.

2) On the periodic table, vertical columns are called ______or families.

3) Is moving in order from Li (#3) to Ne (#10), Na(#11) to Ar(#18) and K(#19) to Kr(#36) represent moving across a period or down a group?

What pattern do you notice as you move across a period from left to right across the periodic table?

4A) As we move from left to right across the periodic table, are the electrons being added to the same valence level or different levels?

4B) As we move from left to right across the periodic table, is the number of positive protons attracting valence electrons increasing or decreasing?

4C) Explain the observed periodic trend.

5A) Are the elements, Li (#3), Na (#11), K (#19) in the same group (vertical column) or same period (horrizontal row)?

5B) As we move down the column, does the size of the atoms increase or decrease?

5C) Write the core electron configuration for each element:

Li: Na: K:

5D) Explain, in terms of electron configuration the observed size trend.

Summary: In the box below, use different size circles to show the size trend across a period moving left to right and down a group. Draw 7 circles down to represent a group and 8 circles across to represent period.

Across a period from L → R, Atomic Radius (AR) ______Reason: Across a period from L → R, (More / fewer) protons are attracting, electrons in (same/different) valence level.

Down a period (top → bottom), AR ______

Reason: Down a period (top → bottom), valence electrons are being placed in higher energy levels farther from the nucleus.

6) For each set, decide whether the elements are in the same period (row) or same family/group (vertical column). Circle the smallest element in each set. (From Zumdahl,p. 336, #82)

A) Na, K, Rb B) Na, Si, S C) N, P, As D) N, O, F

37 http://www.youtube.com/watch?feature=player_embedded&v=uixxJtJPVXk

A question like this that requires you to Link together electron configurations and trends in atomic size to observed physical and chemical properties will be on the test.

1) Compare the Alkali Metals family in terms of physical properties (state of matter, appearance/color, hardness/softness) and chemical properties (do they react with oxygen and water? Do they react at the same rate?) A) What physical and chemical properties are very similar between different elements in the family?

B) Describe how the different elements are different from each other.

Chemical Reactivity and Electron Configurations For each alkali metal, write a core abbreviation electron configuration. Then rank the relative reactivity with oxygen and with water on a scale of 1-5 (1 = Fastest, most violent reaction to 5= slowest, least violent reaction).

Element with core valence e- Reactivity Rank configuration (1-5

Li [He]2s1

2)Fill in the blank: “As we move down the family, the reactions seem to become more______”.

3) Explain, in terms of valence electron configurations, why all of the alkali metals react in a similar manner with oxygen and water.

4) Atoms react in order to obtain filled valence electron levels. Why are the alkali metals all so reactive?

Relative Reactivity, Electron Configuration and Atomic Radius

5a) What were the least reactive and most reactive alkali metals? Least = Most =

5b) Draw circles to show the relative size of each atom. Write the valence electron configuration use core abbreviation under each atom.

6) The data table below shows the relationship between distance between charges measured in nanometers vs Force of Attraction measured in units of Newtons (Data Table from POGIL Activities for High School Chemistry, p. 67, edited by Laura Trout, published by Flinn, 2012).

A) What is the relationship between the distance between + and – charges and the amount of attraction between the charges?

The closer the distance between the two charges the ______the attraction.

6B) Explain why Cs is so much more reactive than Li.

7) Predict which halogen family member would be more reactive, F or I? (Hint #1: Do these atoms want to gain, lose or share e- to fill their outer levels? Hint #2: What are the relative sizes of the two atoms?)

Quantum Unit Review Part 2~ HW 8-13 1) What does it mean to say electron energy levels are quantized?

2) When fireworks, explode they emit different colors of light. The colors are determined by the different elements that make up the firework. For example, Mg releases bright white light while Cu produces a blue-green light. Explain, using a diagram and words, what is happening inside the atoms that causes light to be emitted.

3) What is an orbital?

4) Sketch the following orbitals: i) s ii) p iii) d orbital

5a) Sketch a 1s, 2s and 3s orbital side by side.

Compare: i) electron capacity of orbital (maximum number of electrons that can be held orbital)

ii) Shape of orbital

iii) Size of orbital (average distance from nucleus)

5) For each of the following elements, write complete electrons configurations and orbital notation diagrams (boxes or circles with arrows): i) B, #5

ii) N#7

6) Rank each of the following sets of elements in order of increasing atomic radius (smallest→ largest).

A) Sb, I, Sr B) Te, O, S

7) For each of the following elements, write core abbreviation and orbital notation diagrams (boxes or circles with arrows): A) Si, #14

B) V, #23

C) As, #33

D) Ru, #44

8) The following questions refer to the elements F (element #9), Ne (element #10), Na(element #11) A) Write the complete valence electron configurations and orbital notation diagrams for each element.

Ne:

Na:

B) Which element has an extremely high first ionization energy? Explain.

C) Which element has an extremely low first ionization energy but an extremely high second ionization energy. Explain.

D) Which element is chemically inert (unreactive), which element is extremely reactive and loses 1 electron to form +1 ions, which element is extremely reactive and gains 1 electron to form -1 ions. Explain your answers in terms of the electron configurations you wrote in part a.

E) Arrange the neutral atoms in order from largest to smallest atoms. Briefly explain your answer.

F) Write the complete valence electron configurations and orbital notation diagrams for : i) F-1

ii) Ne

iii) Na+1

Hint: Has an atom with a +1 charge gained or lost an electron? Has an atom with a -1 charge gained or lost an electron? What do you notice about the final electron configurations? Why was an electron gained or lost?

9) Provide an explanation for the following observation: The development of glass panes in windows that allowed natural light into building while blocking wind and cold air dramatically improved comfort inside buildings. Glass typically consists of silicon dioxide with one or more metals mixed in. A common glass made from Si,O and Na was very clear and easy to shape but broke easily. Scientists hypothesized that reason the glass shattered so easily was because Na atoms were too small to make very strong contacts with the network of Si and O atoms. If this were true, then substituting a different atom for Na that the glass might become stronger. Using the periodic table, propose a replacement metal for Na that would have similar chemical properties but a bigger size that could make stronger glass.

10) In Science, good models have predictive power. The Russian chemist Mendeleev is credited with developing the modern periodic table. Mendeleev organized the elements according chemical and physical properties and weight. Mendeleev impressed scientists with his ability to predict the properties of elements that had not yet been discovered. For example, if element 118 were discovered, it would be placed on the periodic table under the element Rn. What phase of matter would you predict element 118 would exist in and what type of chemical reactivity would you predict it would have? Explain.

11) Use the ionization energy data given below to the answer the two questions. The subscript indicates the number of the electron being removed. For example, IE1 is the remove the first electron, IE2 is the energy required to remove the second electron, etc. All values are in kJ/mole.

IE1 = 577 IE2= 1815 IE3 = 2740 IE4 = 11,600 IE5 = 15,000

A) How many valence electrons does this element have?

B) What is the charge on the common, stable ion of this element?

Unit 9 ~ Types of Bonding/Lewis Structure ~ Test Topics

Chemical Bonding • Know the definition of a chemical bond and its importance in determining the chemical and physical properties of a compound (For example, differences in properties between different allotropes(forms) of carbon, graphite, diamond, buckyballs due to differences in bonding patterns). • Understand that chemical attractions are the result of electrostatic attractions between bonding electrons and the positive nuclei of atoms. • Understand that chemical bonds are constantly bending and stretching, and that the best analogy for describing a chemical bond is a spring. Chemists can use the quantized stretching and bending motions of bonded atoms in the infared region to obtain information about the pattern by which atoms are bonded together. Chemical Bonding and Energy • Know that chemical bonds form to lower the energy of a group of atoms. • There are two important factors (driving forces) that lower the energy of bonded atoms: 1) Atoms gain, lose or share electrons to optimize their valence electron configuration. 2) Atoms try to maximize the attractions between opposite electrical charges. • Be able to interpret a graph of potential energy vs. internuclear separation and label bond energy and bond length. • Understand that energy is required to break bonds (separating + and – charges) and energy is released when bonds are formed (+ and – charges coming together). Types of Bonds • Know types of bonding: Ionic, non-polar(pure) covalent, polar covalent and metallic. • Understand detailed bonding models for ionic (formed by transfer of electron(s) from a metal to a nonmetal, nonpolar covalent [equal sharing of electron pair(s)], polar covalent [unequal sharing of electron pair(s)] and metallic bonding (sea of electrons model – mobile, delocalized bonding electrons are shared throughout the entire metal crystal. • Understand how properties such a electrical conductivity be explained by different model of bonding • Be able to identify two different types of covalent bonds: sigma(σ) formed by direct straight on overlap of orbitals and pi (π) bonds formed by overlap of parallel p orbitals. Know that the first bond formed is a sigma bond and if double or triple bonds are present the 2nd or 3rd bonds formed are pi bonds. Electronegativity (EN) • Be able to define electronegativity ( attraction an atom exerts are shared electrons in a chemical bond) and understand two applications of electronegativity differences: 1) EN difference between 2 atoms can be used to predict type of bond (ionic, nonpolar, or polar covalent) 2) EN difference can describe polarity of a bond identify which atom will have full or partial positive or negative charges. Lewis Dot Structures • Be able to draw Lewis Dot structures for both ionic and covalent compounds using appropriate conventions of [ ] for ionic and - for covalent bonds. • Understand when and how to use multiple (double and triple) bonds. • Know exceptions to the octet rule: 1) Electron deficient cases (less than 8 valence electrons) – MUST BE MEMORIZED

BeX2 and BX3 where X = H, F,Cl,Br , I . 2) Expanded octets – recall central atom must be from 3rd row or lower (rows 4,5,6,7) for expanded octets

Notes: Introduction to Chemical Bonds

Chemical Bond: ______of ______that holds atoms together so they act as a ______.

Key Concepts:

1) Chemistry is about ______charge attracting ______charge. All chemical bonds, regardless of type are the result of

In a chemical bond, the bonding electrons are simultaneously attracted to the positive nuclei of the two atoms.

2) The type of chemical bonding in a substance

Example #1: 3 Forms of Carbon: Graphite, Diamond, Buckyballs have different properties due to differences in

Example #2: Salt (______Bonding), Sugar (______Bonding), Iron (______Bonding)

3) Making and Breaking Chemical Bond always involves ______. Breaking chemical bonds costs energy (energy required to separate + and – charges), forming chemical bonds releases ( - charged electrons are spending more time near + charges).

Magnet Analogy: Pulling two magnets apart requires work but the magnets will move together on their own when placed nearby.

Net energy change = difference between cost of breaking old chemical bonds and energy released when new bonds are formed.

Question: Why do chemical bonds form?

Answer: The formation of chemical bonds ______of a group of atoms. Atoms bonded together are more ______than separated individual atoms.

Lower Energy = ______stable! Higher Energy = ______stable!

Two Reasons Chemical Bonds Form:

1) Individual atoms seek to ______their ______configurations by ______, ______, or ______valence electrons.

3s23p5 2s22p6 3s23p6 3s1

2) Bonding electrons are ______attracted to ______nuclei. (When – electrons spend more time near + charges this lowers the energy of the atoms = more stable).

Bond

Lower E =

4) Chemical bonds are generally formed by the ______of ______. Diagram of two 1s orbitals from H overlapping to form molecular orbital of H2.

Two different types of Covalent Bond orbital overlap: Sigma and Pi Sigma (σ): first bond formed between 2 atoms is the result of “ ______-____” overlap of 2 atomic orbitals.

Diagram of s-s (σ): overlap

Diagram of s-p (σ): overlap

Diagram of p-p (σ): overlap

Pi (π) – 2nd and 3rd pairs of shared are bonded by “______” overlap of parallel ____ orbitals.

Diagram of p-p (π) overlap:

Single Bond share 1 pair of e-: ____ σ bond and _____π bond(s)

Double Bond share 1 pair of e-: ____ σ bond and _____π bond(s)

Triple Bond share 1 pair of e-: ____ σ bond and _____π bond(s)

5) Chemical Bonds are not static (stationary); bonds can ______, ______or in some cases twist or rotate.

The best analogy for a chemical bond is a ______

. Molecules in bonds bend, twist or in some cases rotate when they absorb ______light.

Different bonding patterns (different atoms and single vs double vs triple bonds) absorb IR light at different wavelengths and hence different quantized energy levels. The absorbance band wavelength and appearance can be used to determine the type of bond.

Ionic, Nonpolar Covalent, Polar Covalent and Metallic Bonding

Classifying Bonds: The classification of ionic and covalent bonds is based on the extent to which electrons are ______or ______.

Complete sharing Partial Sharing, Partial Transfer Complete Transfer

______

ELECTRONEGATIVITY

Defn-

Pauling Scale: Ranges from 0.0 (no attraction) to 4.0 (Highest Attraction); Label most electronegativity, least electronegative areas; Which family is missing from chart below and why?

Connection between atomic radius and electronegativity:

Smaller atoms generally have electronegativity.

IONIC BONDS

A complete ______of 1 or more ______from a ______to a ______forming ____ and _____ ions.

Example: Sodium ______to chlorine to form sodium chloride.

Na + Cl → Na Cl

Trillions of individual ions bond together to form a “______”.

Note: Central - charged chloride is surrounded by + charged sodium ions

NON-POLAR COVALENT BONDS

An ______(co) of ______electrons. (“valent”).

Picture of Cl2: Molecule:

Typically forms between ______or two nonmetals with a very similar attraction for electrons. (Similar electronegativity values). Examples: H or Cl 2 2

POLAR COVALENT BONDS

An ______of ______electrons.

Picture of HCl Molecule:

Typically forms between two ______nonmetals.

The ______will be “______” ______and have ______

______around it. Arrow points towards ______side of molecule.

Picture of HF including partial charges ( ) and arrows:

Summary of Bond Types between Atoms

Where M – metal, N = nonmetal; subscripts indicate same or different

Type of Bond Element Bonding Pattern Description Example

Metallic Bonding – “Sea of Electrons” Model

• In Metal Crystal, electrons are ______and ______(constantly moving throughout the entire metal crystal.)

• Bonding in metals can be visualized as

Diagram of “ Electron Sea Model”:

= + cations (nucleus + core e-)

e- = valence “delocalized” electrons

Use to show motion of electrons Properties of Metals Explained by Sea of Electrons Model

1) ______– metals appear ______white because they ______essentially _____ wavelengths of light. Explanation: Electrons constantly moving in different positions near many different nuclei experience a hugh range of energies and hence can make nearly every possible quantum energy transition possible which releases light in the visible range of light.

2) ______- metals can be easily ______(often this requires metals at high temperatures, reshaping them and then allowing them to cool. Some metals can easily bend without breaking at room temperature.

3) ______- metals can be easily drawn into a ______. Example ___ gram of gold can be drawn into a wire _____ mile long.

Explanation for Observation for Metals being Ductile and Malleable:

In covalent and ionic bonds, the positions of + nuclei and – electrons are fixed relative to each other, so if the nuclei shift positions the bonds will break.

In the case of ionic bonds, shift of nuclei can cause like charge ions (+ and + or – and - ) to come too close. The repulsion pushes the crystal apart.

In the case of covalent bonds, electron must stay directly between nuclei or bond will break.

Since electrons in a metal crystal are delocalized, metal electrons can shift their positions to still be attracted to different nuclei within the crystal when the nuclei shift positions.

4) ______conduction Explanation: ______, ______, ______rapidly carry the kinetic energy to different atoms throughout the metal crystal

5) ______Conduction: Explanation: Electrical Current = ______of ______. Mobile (______) ______carry current in metals. Electrons ______metal crystal at negative electrode and ______at the positive electrode.

# of electrons entering # of electrons exiting

+ electrode (e- exit) electrode (e- enter)

6) Metal alloys – pure metals can be made harder and stronger by bonding with another metal or with carbon (carbon + iron = steel).

IONIC LEWIS DOT STRUCTURES

Lewis Dot – notation for showing valence electrons and bonds within a compound

Use to represent valence electrons; Use [ ] to represent ionic bonds; charge written in top right corner.

Dot of elements (Recall # valence e- = group #)

Group 1 Group 2 Group 3 Group 4 Group 5 Group 6 Group 7 Group 8

Li Be B C N O F Ne

Draw an ionic Lewis Dot Structure for each of the following: Use [ ] for bonds and alternate + and – charges.

NaCl: Na Cl → [Na]+1[ Cl ] -1

MgCl2 Cl Mg Cl → [ Cl ] [ Mg] [ Cl ]

Li2O

AlBr3

Covalent Lewis Dot Structures Developed by American chemist, G.N. Lewis, Lewis Dot Structures are a Notation system for representing shared (covalently bonded) as lines and nonbonding valence electrons as dots in a molecule.

= Nonbonding electrons (Lone pairs) assigned to one atom. Each dot represents ___ Electron. = single bond (shared pair): Each dash represents 1 pair or ____ electrons.

Example: H ∙ + ∙ H → H∙∙ H = H – H

= double bond = _____ shared pairs or ______total electrons. = triple bond = _____ shared pairs or ______total electrons.

Why do bonds form? Atoms gain, lose or SHARE valence electron to obtain a ______valence electron level.

Final structures (with a few specific exceptions discussed later) must always have a filled valance level for each atom:

DUET RULE: Applies only to ______. Each hydrogen atom will make ____ bond in order to completely fill the ______sublevel.

OCTET RULE: All other atoms besides H (+ octet rule exceptions discussed later) will always total between shared and nonbonding electrons ______valence electrons to obtain a ______valence electron configuration.

In the structure below: Indicate the number of bonding, nonbonding and total valence electrons for each element.

Element Bonding electrons Nonbonding Electrons Total Valence Electrons H Cl C

Practice applying the duet and octet rules: For each molecule add in the number of dots necessary to fill the valence level.

1) H – F 2) 3)

4) This structure has no dots, but is missing 1 or more bonds. Add in the missing bond(s) to provide each atom with a filled outer energy level.

STEPS for Drawing Simple Lewis Dot Structures 1) Draw Skeleton Structure (see notes below on determining central atom), making sure each atom is forming its expected number of bonds. 2) Distribute Dots so each atom has filled outer level

How to Determine the Central Atom in a Lewis Dot Structure Most of our Lewis Structures will have a pattern will a central atom with attached atoms on the outside.

Examples:

1) Method #1: Central Atom = Single Atom in Formula

AX2 , AX3, AX4, AX5, AX6 : A = single atom, X = different atom type

Examples: What is central atom in the following examples?

CO2 : NF3: CCl4: PCl5: SF6

2) Method #2: Central Atom = atom that can make the greatest # of bonds # of bonds = # of electrons needed to fill outer energy level Example: HOF

Element # of valence e # of e’s needed to fill # of bonds formed outer level

H 1 1 1

O 6 2 2

F 7 1 1

Element that forms greatest # of bonds = ; Central Atom would be

Drawing Lewis Structures: Quick Method for Simple structures: ➢ Central Atom: ➢ Draw skeleton- make sure each atom has appropriate number of bonds it typically makes (from HW 9-7). ➢ Add dots to ensure each atom has filled outer level (H = duet, all other atoms = octet)

Example #1: HOF Example #2: CH2O

Drawing Lewis Structures: Electron Count Method for octet rule exceptions and complex problems. Example #1: CH2Cl2 using count method

1) Step #1: Determine Central Atom and connect all atoms with single bonds.

Steps #2-4: Determining total # of Dots available for structure: 2) Count Total Number of valence electrons from each atom in molecule. (Recall: Group # = # of valence e- ) C: H2: Cl2:

3) Count # of bonds and multiple by 2 to obtain total # of bonding electrons. Subtract # bonding electrons from total valence electrons to determine # of dots. Total valence e- - Bonding e- (# of bonds x 2) = # of dots

Total valence e-: - Bonding e-: # of dots:

4) Distribute Dots so each atom has an octet. Notes: 1) Give outside atoms and most electronegative atom octets first. Any leftover dots have to go on central atom.

2) Double and triple bonds - are required when central atom does not have an octet and no more dots are available; Have’s (atom with extra dots) shares a pair with have nots (atoms needing) to form additional bonds. Example: CS2

Important Note on Electron Counting: If a molecule has an overall + charge, add 1 electron to the total count for each +; If a molecule has an overall - charge, add 1 electron to the total count for each +; +1 Example #1: OH- Example #2: NH4

O: 1 atom x 6e- per atom = 6 N: 1 atom x 5e-/atom = 5

H: 1 atom x 1e-atom =1 H: 4 atoms x 1 e-/atom = 4

7 + 1 (from -1 charge) = 8 total valence e- 9 - 1 (from +1 charge) = 8 total valence e-

OCTET RULE EXCEPTIONS:

1) ELECTRON DEFICIENT - Memorize BeX2 and BX3 where X = H, F, Cl, Br, I

Example #1: BeCl2 - NOT IONIC (although metal + nonmetal)

- NO DOUBLE BONDS (although Be does not have octet; halogen never double bond).

Be is stable with only _____ valence electrons

Example #2: BF3 - NOT IONIC (although electronegativity difference = 2.0 > 1.7 predicts ionic [ B = 2.0, F=4.0]) - NO DOUBLE BONDS (although Be does not have octet; halogen never double bond). - B is stable with only _____ valence electrons.

2) EXPANDED OCTETS – central atom from periods _____ thru ______, may be stable with _____ than ______valence electrons.

-1 Example #1: PCl5 Example #2: IF4

Unit 9 Bonding Assignment List

9-1 Race to Catch to a Buckyball

9-2 Sec 12.1,12.2,12.4,12.5 Chemical Bonding Problem Set (wb)

(p.341 – 346, 347 – 352)

9-3 Double, Triple, Sigma and Pi Bonds

9-4 Types of Chemical Bonding (workbook)

9-5 Pages 441, 442 Metallic Bonding (workbook)

9-6 p. 437-440 Properties of Different Bond Types (electrical conductivity)

9-7 Lewis Dot Structure Patterns

9-8 Sec. 12.6, 12.7 (p. 352 – 363) Lewis Dot Structure Problem Set

9-9 Lewis Dot Exceptions and Review

9-10 Unit 9 Review Problem Set

HW 9-1 Race to Catch a Buckyball

This PBS NOVA episode describes the discovery in the last part of the 20th century of a brand new form of the element carbon. Prior to this discovery, it was assumed that there were only two forms of elemental carbon, diamond and graphite. This video illustrates the twists and turns in making a Nobel Prize winning discovery. It also serves an excellent illustration of most important concepts in chemistry – how atoms bond together determines the properties of a substance. Cast of Key Characters

Wolf Kratschmer (Heidelberg) and Donald Huffman (University of Arizona) – in early experiments with vaporizing graphite sticks under low pressure found a mysterious unknown material that had a mysterious “double camel” hump UV spectrum that didn’t match any known substance. At the time, they didn’t realize the “junk” they had made was actually a new form of carbon and thus missed out on the Nobel Prize.

Rick Smalley (Rice University) and Harry Kroto (University of Sussex) shared the Nobel prize with Bob Curl(not featured in video) for the realization that the carbon could bond together in closed cage structures they named Buckminister Fullerenes or Buckyballs for short (named for an architect famous for building geodesic domes).

1) Stars are believed to form when giant clouds of hydrogen are attracted together by gravity. The incredible strength of gravitational attraction from these massive cloud pulls the hydrogen atoms together so violently the nuclei of the atoms bond together in a process called nucleus fusion. What does Kroto mean when he says “we are all made of star-dust”?

2) The story actually starts not with the search for a new form of carbon, but rather trying to understand the origin of mysterious interstellar absorption bands. In our unit on Quantum mechanics we discussed how emission spectra are produced. What is happening inside the atom that causes specific energies of light to be absorbed? Absorption spectra are referred to as “fingerprints” for an element or a compound. What does this mean?

3) Carbon is known to bond to itself or with hydrogen to form long chain molecules. Harry Kroto hypothesized that the mysterious “interstellar absorption bands” might actually be caused by clusters of long chain carbon molecules. He traveled to Rick Smalley’s lab at Rice University in Houston, Texas. Smalley’s lab studied how silicon atoms bonded to together in a high temperature laser in a vacuum. A) Smalley’s group was studying silicon chemistry in the mid-1980’s because of a new technology called the personal computer. What does silicon have to do with computers?

B) Use your knowledge of the periodic table to explain why a laser system designed to study silicon could be useful for studying carbon chemistry.

C) How are the conditions of the laser beam similar to conditions in deep space?

4) In Kroto and Smalley’s experiment carbon in the form of graphite was vaporized in a high temperature laser and the results were analyzed using a mass spectrometer. Recall from first semester that mass spectrometers work by first ionizing (giving an electrical charge) molecules and then measuring the amount of deflection each molecule experiences as it passes by an electromagnet. The lighter molecules deflect more, heavier molecules deflect less.

The results of experiment were extremely surprising. The height of each peak represents the relative abundance of the peak and the number on the x-axis indicates the number of carbon atoms. The researchers expected to see all different sizes of carbon molecules in similar abundance. The actual results are shown above. Which peak is by far the most abundant?

5) Smalley and Kroto used molecular models to try to understand what structure with 60 carbons could be so stable. They started by looking at the models for how carbon bonds.

A) What are the two forms of carbon called?

B) Describe the physical properties and appearance of each form of carbon.

C) Which form of carbon bonds together with 4 other carbon atoms and has hydrogen atoms on its outer surface?

D) Which form of carbon bonds to 3 other carbon atoms and bonds together in chicken-wire like sheets?

The forces of attraction within a layer are very ______, while the forces of attraction between layers are very ______.

E) Using the fact that carbon atoms need to form 4 different covalent bonds with other atoms and that only certain angles are possible between bonded atoms, Kroto and Smalley proposed a novel cage-like structure that looks like a ______ball. It contains 12 ______surrounded by 20 ______. They named their new structure ______after the architect Buckminster Fuller who was a pioneer in the architecture of geodesic domes.

F) “The Long Ranger and Tonto”. Kroto noticed that the big peak was accompanied by a smaller peak at 70. Which picture below matches the 60 carbon peak and which picture matches the 70 carbon peak?

Nanotubes

An important application of buckyballs is to construct very strong lightweight fibers called nanotubes. Buckyballs and nanotubes are incredibly strong and are excellent electrical conductors.

6) Role of Theory in Practical work: Kroto and Smalley’s interpretation of their data was criticized because of their inability to synthesize compound in large enough amounts to hold in one’s hand and be able to carry out other experiments to verify the structure. Because of the symmetry of the C60, the ______spectrum (due to electrons absorbing ultraviolet light) and the ______spectrum (due to bending and stretching of bonds from absorbing infrared light) were able to be predicted. When these predicted spectra calculated from theory were published, Wolfgang Kratschmer (Heidelberg) and Donald Huffman (University of Arizona) recognized these predicted spectra as a match for the “junk” spectrum they had made years before. Their apparatus for making C-60 involves running high voltage electricity through graphite under a low pressure helium atmosphere.

Why is helium used in this experiment instead of air? (Hints: How is helium chemically different than oxygen in air?

HW 9-2 Introduction to Chemical Bonding Problem Set

1) Chemistry is about __ charge attracting __ charge. All bonds are the result of attractions between ______electrical charges.

2) Bonding, Chemical and Physical Properties A) The English language contains at least 250,000 distinct words but only 26 letters. How is it possible to make so many words from such a small number of letters?

B) The number of distinct chemical substances in the universe is unknown but it is estimated to be between 1018 and 10200. The number of naturally occurring elements is approximately 100 or 102. How is it possible to make so many different substances from such a relatively small number of different types of atoms?

C) Diamond, graphite and buckyballs are all made of carbon and yet they have very different chemical and physical properties. For example, graphite is more chemically reactive than diamond or buckyballs. Diamond is much harder than graphite and buckyball are very resistant to being broken down by light. How can we explain these different properties if they are composed of the same substance?

3) What would the universe be like if chemical bonds did not exist? (Hint: Think about objects around you that contain chemical bonds)

Energy changes and Chemical Bonding 4) All chemical reactions involve a change in ______, due to the breaking and making of chemical bonds. In chemical reactions old bonds break and new bonds form. Breaking bonds cost or requires energy while forming bonds releases energy. The overall energy change for a reaction is the net difference between cost of breaking bonds and the energy released when new bonds are formed.

5) List two reasons chemical bonds form or driving forces (factors that lower the energy of the atoms and hence make them more stable) for the formation of chemical bonds.

i) ______valence electron configuration by ______, ______, or ______electrons.

ii) Bonding electrons are ______attracted to two positive nuclei. (Lowers energy by allowing – charge to spend more time near + charge).

6) Consider the diagram below showing interactions between the protons of 2 different atoms, the electrons of two different atoms and the attractions of protons for electrons. Label each type of interaction as attractive or repulsive.

i) Electron (-) from one atom with Electron from different atom ______ii) Protons (+) from one atom with Protons (+) from different atom ______iii) Electron(-) from one atom with protons from the nucleus of the same atom. ______iv) Electron (-) from one with protons from the nucleus of the other atom. ______

7) The diagram below shows the energy changes that take place as 2 H atoms approach each other.

A) What is the H-H bond length (include appropriate units)?

B) What is the energy change for the formation of an H-H bond (include appropriate units)?

C) What is the energy change for the breaking of an H-H bond (include appropriate units)?

For questions, D-G, Match the picture #4,3,2,1 depicting the interaction of the 2 H atoms with the description: D) At this point the atoms are at their lowest energy point, where attractive forces between + and – charges are at a maximum compared to repulsive forces between like charges.

E) At this point, the atoms are too close together and the nuclei are pushing the atoms away from each other. A hugh amount of energy would be required to push the atoms any closer together.

F) At this point, the atoms are too far apart for the electrons from one atom to feel that attraction for the other atom.

G) At this point, the atoms are beginning to approach each other closely enough for each atom’s electron to start to experience attraction for the nucleus of the other atom.

HW 9-3 Single, Double, Triple Bonds – sigma and pi bonding overlaps of atomic orbitals

1) Rank the following pictures in order of INCREASING attraction between + and – charges. (From POGIL, Activities for High School Chemistry, edited by Laura Trout, Flinn Scientific Publishing, p.73.)

- + + - +

+ -

2) Consider a C-C single bond, a C=C double bond and a C C

A) Match each picture with the appropriate bond (Each C has the same amount of + charge, +6, in its nucleus; shared electrons are represented as e-.)

e- + + e-

e- e- + + e- e-

e- e-

+ e- e- +

e- e-

B) Predict which bond would be the strongest. Explain your prediction in terms of the electrical charges involved. C) Predict which bond length would be the shortest. Explain your prediction in terms of the electrical charges involved. 63

3) The graphs below represent a C-C single bond and a C=C bond. Which graph represents which type of bond? Explain your reasoning in terms of bond energy and bond length. (From POGIL, Activities for High School Chemistry, edited by Laura Trout, Flinn Scientific Publishing, 2012, p.230.) .

4) Label each picture as a sigma bond (σ) or a pi bond (π)

5) How many sigma and how many pi bonds are present in each of the following molecules?

A) CH2O σ: π: B) HCN σ: π:

B) CO2 σ: π: D) C2H2 σ: π:

6) Explain the following observation:

The average strength of a C-C sigma bond is 347 kJ/mol while the average strength of a C-C pi bond is only 267 kJ/mol. (Hint: Use the diagram below to study the average distance the e- [represented by arrows]in each type of bond are from the positive nuclear charges).

7)Multiple Choice: Choose the letter of the best answer. A sample of CH2O gas and a sample of HCN gas are each tested by shining infrared light through each sample. The pattern of quantized energy absorbance for each sample is different because the stretching and bending motion of a C-O double bond is different than the stretching and bending motions of a C-N triple bond. The stretching and bonding motions of the bonds that explain the different quantized spectrum can best be understood by visualizing chemical bonds as behaving like

A) A wooden stick B) a metal spring C) a plastic stick

Types of Chemical Bonding Problem Set ~ Homework 9-4 For questions #1-12, use the letters I, NPC, PC or M to describe the type of bonding. I for Ionic NPC for Non-polar Covalent PC for Polar Covalent M for Metallic

1) Unequal sharing of e- between 2 atoms 2) Equal sharing of e- between 2 atoms

3) Delocalized of Sharing of e- between all atoms 4) Transfer of one or more e- to form + and – ions

5) Type of bond that forms between Li + O atoms 6) Type of bond that forms between Cu atoms (Cu + Cu)

7) Type of bond that forms between C + O atoms 8) Type of bond that forms between N + H atoms

9) Type of bond that forms between two F atoms (F + F) 10) Type of bond that forms between H + O atoms

11) Type of bond that forms between Cu + Cl atoms

12) Label each of the picture below as representing in ionic, nonpolar covalent or polar covalent.

a) b) c) Metallic

13) Define the term electronegativity:

14) Given the following pairs of atoms with the electronegativity values written next to each atom:

a) 2.1 H – O 3.5 b) 2.1 H – Si 1.8

• Label atoms in the bonds above that are δ- . (Higher electronegativity value) • Label atoms in the bonds above that are δ+ . (Lower electronegativity value) • Draw an arrow below each molecule with the arrow head pointing towards the partial negative end. • Which bond, the H-O or the H-Si is more polar (more polar = greater EN difference)? Metallic Bonding ~ Homework 9-5 1) List 5 properties of metals.

2) Use the properties of metals that you listed in question #1 to help answer the following. A) Outline the properties of bronze (an alloy of copper and tin) that make it better material for making weapons and tools than stone.

B) What two properties of metals such a copper, have been critical in our ability to construct a modern power grid in which electricity is delivered directly to homes and workplaces?

3a) Using a diagram (label positive ions and use arrows to show moving electrons) and words, describe the “electron sea” model of metallic bonding.

3b) What does it mean to say that electrons are mobile and delocalized?

4) Electric current = ______of ______. How does the model of metallic bonding explain why metals are excellent conductors of electricity?

5) Pure metals are typically very soft. Alloys are formed by mixing two metals (example copper + tin = bronze or iron and carbon + steel.) How does the amount of the carbon relative to the amount of steel change the properties of steel?

Homework 9-6 – Electrical Conductivity and Bonding 1a) Fill in the chart below: Type of Bond Conducts as solid? Conducts as liquid? Conducts as aqueous solution?

Ionic

NP or Polar Covalent

Metallic NA – not water soluble

1b) Under conditions in which an ionic substance conducts electricity, what is the particle that carries the charge?

1c) Under conditions in which a metallic substance conducts electricity, what is the particle that carries the charge?

3) Explain the following observations by identifying the type of bonding present and whether moving charged particles are present. a) Pure, distilled water is a poor conductor of electricity and solid salt is a poor conductor of electricity but salt water is an excellent conductor of electricity.

b) Pure copper (Cu) is an excellent conductor of electricity, but CuCl2 is a poor conductor of electricity.

HW 9-7 Covalent Lewis Dot Structure Patterns

Use lines and dots to show how each atom typically forms bonds in Lewis structures.

Hydrogen

# of Valence e- (Group #) # of electrons # of bonds formed Common bonding pattern: needed to fill outer energy H level H always obeys the ______H: rule. (duet or octet)

Halogen Family, Group 7 (Family 17)

# of Valence e- (Group #) # of electrons # of bonds formed Common bonding patterns: needed to fill outer energy F Cl level

F: Br I Cl:

Br: Halogens obey the (duet or octet)

I: ______rule.

Group 6 (Family 16)

# of Valence e- (Group #) # of electrons # of bonds formed Common bonding patterns: needed to fill outer energy O or O level O: S or S S: Se: Se or Se

Group 6 obeys the (duet or octet) ______rule. 68

Group #5 (Family 15)

# of Valence e- (Group #) # of electrons # of bonds formed Common bonding patterns: needed to fill outer Note: N makes strong pi bonds and hence energy level will form double or triple bonds, while P N: and As typically only form single bonds.

As:

N or N or N P As

Group 5 (Family 15) obeys the (duet or octet) ______rule.

Group #4 (Family 14) # of Valence e- (Group #) # of electrons # of bonds formed Common bonding patterns: needed to fill outer energy Note: C makes strong pi bonds and hence level will form double or triple bonds, while Si typically only forms single bonds. C:

Si:

C or C or C or C Si

Group 4 (Family 14) obeys the (duet or octet) ______rule.

Conclusion Question: What is the relationship between number of bonds formed and number of electrons needed to fill the outer energy level?

Multiple choice: For each pattern, identify the element that could represent X. Choices: A) H B) F C) C D) N E) O

1) X 2) X 3) X 4) X 5) X

Homework 9-8 Lewis Dot Structures-

Part 1: Ionic Lewis Dot Structures (Metal + Nonmetal) – Draw Lewis Structures for the following ionic compounds. Use [ ] to represent ionic bonds with charges written in the top right hand corner.

A) KBr B) CaCl2 C) Al2O3

Parts 2 -6 are covalent structures (Nonmetal + Nonmetal).

Part 2: Single Covalent bonds, no formal charges

A) NH3 B) H2O

C) HOCl D) CCl4

Part 3:( Single bonds with overall charges)

-1 +1 A) OH B) NH4

Part 4: Double Bonds

A) CO2 B) CH2O

Part 5: Triple bonds

A) N2 B) C2H2

HW 9-9 Octet Rule Exceptions and Review

Part 1: Octet Rule Exceptions

A) BeF2 B) BCl3

C) BeBr2 D) BBr3

-1 E) XeF4 F) I3

Part 2: Test Format questions. For each compound, first indicate the type of bond that will form ionic (metal + nonmetal) [ ] or covalent - (Nonpolar or polar is same notation- two nonmetals) and then draw the correct Lewis structure.

A) H2Se B) CaF2 C) BeH2

D) N2Br2 E) CO F) MgO

-2 G) NF3 H) SO4 I) XeF2

Bonding Unit Review ~Homework 9-10

Use the pictures below to answer questions #1-4

A B C 1) Which structure represents graphite? 2) Which structure represents diamond? 3) Which structure represents a buckyball? 4) Explain why graphite, diamond, and buckyballs have such different properties if they are all made of carbon?

For Questions #5-8, evaluate each statement as true or false.

5) Chemical bonds between two atoms are very rigid. A good analogy for bond is a steel rod. 6) One driving force for the formation of chemical bonds is to increase the energy of the group of atoms. 7) One driving force for the formation of chemical bonds is to enable atoms to obtain a completely filled valence electron energy level. 8) The force that holds atoms together is electrons being simultaneously attracted to two positively charged nuclei. Use the diagram below to answer questions #9 -11.

9) Which point represents the lowest energy, most stable point? 10) Which point represents the two atoms too far apart to attract each other? 11) Label the bond length and bond energy on the graph.

12) Given the following electronegativity values: S: 2.6 N: 3.0 O: 3.5 A) For each bond, indicate which atom will have the partial positive and which atom is partial negative.

i) S – O ii) N – O

B) Which bond is more polar?

13) Bond Energy : Fill in the blank: Breaking bonds ______energy (pushing + and – charges apart) while making bonds ______energy (+ and – charges pulled together).

Use the diagrams below to answer questions #14-16

A) B) C) 14) Which diagram represents a polar covalent bond? 15) Which diagram represents an ionic bond? 16) Which diagram represents a nonpolar covalent bond?

17) While you are up late at a coffee shop studying chemistry with your friends, you suspect that one of your “friends” has switched the sugar and salt dispensers so that you will put salt in your coffee instead of sugar. For a person not trained in chemistry, this practical joke has a chance of success because of course both salt and sugar are white crystalline solids. You however, pull out your portable crucible, Bunsen burner and electrical conductivity sets and record the following data: Test Compound X Compound Y

Conductivity of solid NO NO Conductivity of liquid NO YES Conductivity of aqueous solution NO YES

Which set of results, Compound X or Compound Y is consistent with sugar? Explain in terms of the bonding involved for salt and sugar. If X represents the material in the sugar dispenser, did your friend actually switch the sugar and salt?

Use the table below to answer questions #18-20: Table 1 Properties of Mystery Compounds Substance I Substance II Substance III Melting Point 419°C -196°C 680°C Electrolyte in water No No Yes Luster High Low Low Electrical Conductivity (in solid state) High Low Low Thermal Conductivity High Low Low Malleable and Ductile Yes No No Density 7.1 g//cm3 0.00125 g/cm3 3.54 g/cm3 Use your knowledge of the type of bonding present in each substance to identify the substance based on its physical properties. Briefly explain your reasoning.

Matching Choices for 18-20: A) ZnSO4 B) Zn C) N2

18) Substance I = 19) Substance II = 20) Substance III =

Use the diagrams below to answer questions # 21 and 22

A B)

21) Which diagram could be a model of the bonding in sodium metal atoms?

22) Which diagram could be a model of the bonding in sodium chloride?

23) Explain why sodium metal conducts electricity while solid sodium chloride does not.

24) Explain why molten (liquid) sodium chloride or sodium chloride dissolved in water are excellent conductors of electricity.

LEWIS DOT- Identify the bond type (ionic or non/polar covalent) and Draw correct Lewis Dot structures for each of the following compounds. Be sure to use [ ] to represent ionic bonds with correct charges and lines to represent covalent bonds and dots to represent nonbonding electrons.

1) MgF2 2) PCl3 3) BF3

4) BeF2 5) KBr 6) N2F2

-3 7) HOBr 8) C2F2 9) PO4

+1 10) HCOOH 11) IF4

UNIT 10 ~ VSEPR/Dipoles/ IMF’s ~ Test Topics

Valence Shell Electron Pair Repulsion Theory • Understand the basic principles of VSEPR, including the idea that valence electron pairs around the central atom of a molecule will arrange themselves as far apart in space as possible in order to minimize electrostatic repulsion. • Given the chart of geometry names (based on the number of bonding and nonbonding electron pairs around the central atom), be able to predict the geometry of a compound. • Be able to draw three-dimensional representations of the geometries listed on your VSEPR handout and Molecular Models Lab. • Know the bond angles for geometries for base shapes (linear, trigonal pyramid, tetrahedral, trigonal bipyramid, and octahedral.) • Remember that the 3,1 trigonal pyramid and 2,2 Bent structures are “cousins” of the 4,0 tetrahedral. Know that nonbonding pairs exert a greater repulsion and hence compress bond angles more than bonding pairs. (DO NOT have to memorize 107 and 104.5 angles but should be able to rank tetrahedral, trigonal pyramid and bent angles in terms of relative sizes.

Dipole Moments • Be able to determine whether or not a molecule has a net dipole moment. If a molecule does have a net dipole moment, indicate the direction of the dipole moment. • For example, water has a bent or angular shape due to the two nonbonding(lone) pairs on the central oxygen. The bond dipoles point toward the more electronegative O atom.

This creates a NET or overall dipole moment that points straight “up” from the H atoms towards the top of the O atom. This net dipole makes water a POLAR molecule.

Intermolecular Forces

• Know the 3 basic types of weak intermolecular forces: dipole-dipole, hydrogen bonding and London dispersion (LDF). Be able to predict the predominant type of attraction that will occur between two molecules. • Understand the typical relative strengths of the 3 forces (H-bond> dipole> LDF) and be able to use this information to predict relative melting and boiling points of substances. Understand that strength of LDF depends on size of electron cloud. .

Valence Shell Electron Pair Repulsion (VSEPR) Theory

Model for predicting three –dimensional shapes (geometries) of molecules.

Importance: Our universe is three-dimensional – true properties of molecules can only be understood by looking at actual arrangements of molecules in space. Example application: Chiral molecules – most drugs contain a carbon bonded to 4 different other atoms or groups of atoms. In a flat 2-D picture in appears arrangement of atoms does not matter; In reality, different arrangements of molecules are mirror images of each other and your body will only accept one form and rejects the other.

Key ideas: Molecular Geometry is determined by

1) The arrangement of ______and ______valence electron ______around the ______atom in a molecule.

2) The bonding and nonbonding valence electron pairs arrange themselves as far ______in space as possible in order to minimize ______between negative charges.

Key points for counting Bonding and Nonbonding electron pairs on CENTRAL ATOM: 1) Double and triple bonds count as only ONE bonding pair (one bonding “region”). 2) Count Nonbonding electron pairs around CENTRAL ATOM ONLY. Represented as

Nonbonding electrons repel more than bonding pairs. (Need to be able to rank bond angles from largest to smallest, don’t have to memorize exact numbers for test trigonal pyramid and bent but need correct numbers for lab).

Example: Bond Angles: CH4 = NH3 = H2O =

Conventions for drawing in 3-D:

Use to represent bonds in plane of paper

Use to represent bonds receding into plane of paper

Use to represent bonds coming out of plane of paper

Use to represent nonbonding (lone) electron pairs on central atom

Notes ~ Dipoles and Polar Molecules Dipole (“di” means 2, dipolar means two poles) • An ______distribution of electron density in a molecule. • Results in a partial _____ (δ + ) and partial ______(δ - ) sides of molecules. • The direction of the dipole is shown as an ______pointing towards the ______

end (the cross on the arrow is near the ______end).

H - F

Polar Molecules - Molecules with a “net” or overall dipole. 3 Steps for determining polarity Step #1:

Example: H2O

Step #2: Bonding Regions? ______Nonbonding regions? ______

Geometry of water = ______

Step #3: 1) Are ______Bonds Present? (Polar bond = ∆EN, an electronegativity difference between bonded atoms.) If NO Polar Bonds are present molecule will be ______. Example:

2) Nonbonding Pair on ______atom? (Most important when outside atoms are H atoms which lack nonbonding pairs)

3) ______of molecule– Are polar bonds and nonbonding electrons pairs arranged symmetrically?

NO – NOT Symmetric YES - Symmetric

If symmetric = ____ NET DIPOLE (NONPOLAR) NOT symmetric = NET DIPOLE (POLAR)

Example: CO 2 Example: H2O Polar Bond dipoles cancel; equal and opposite Polar Bond dipoles reinforce 77

Notes: Intermolecular Forces (IMF’s)

Intramolecular – forces of attraction ______molecules

3 types of intramolecular forces:

______(Ex: NaCl) ______(Ex: H2, HCl) ______(Ex: Cu, brass)

Intermolecular – forces of attraction ______molecules. Much weaker than chemical bonds ______molecules.

The strength of IMF’s is linked to phase of matter. Stronger IMF’s = ______mp and bp.

“Biscuit” (Cookie) Analogy: The icing joining two cookies is an analogy for a ______bond.

The analogy for intermolecular forces between biscuit molecules is ______.

Solid line = Bond (strong) 3 Types of Intermolecular Forces

Dotted line = Force (weak)

Order of magnitude approximation of “typical” relative strengths: (Note: Actual values for type vary over a range) Covalent bond = 100 units Hydrogen bond = units Dipole-Dipole = units LDF = units

Dipole- Dipole – weak attractive force between ______molecules. The partial _____ end of one molecule is attracted to the partial ______end of another molecule.

Example: Dipole-dipole between IF molecules

Hydrogen Bonding: An especially strong type of dipole-dipole attraction between molecules.

Requirements for hydrogen bonding: H-FON and FON

1) ____ covalently bonded to _____, ______, and ______, AND

2) ______, ______, ______, present on second molecule (N must have available nonbonding pr)

Examples of Hydrogen Bonding:

H2O (show 4 other water molecules) NH3 (show 1 other NH3) HF ( show 2 other HF)

Practice: Which pairs of molecules can form hydrogen bonds? Explain.

A) H2O and NH3 B) CH2O and CH2O

Why is a hydrogen bond stronger than a typical dipole-dipole attraction? Strength of attraction between + and – charges depends upon:

1) magnitude of charge (i.e +1 and -1 is weaker than +3 and -3) 2) distance (closer = stronger) Electronegativity values: H = 2.1 F = 4.0 O = 3.5 N = 3.0

Large EN differences: F- H → 4.0 -2.1 = 1.9, O-H → 3.5 – 2.1 = 1.4 N-H → 3.0 – 2.1 = 0.9 means very polar bonds and therefore large partial positive and partial negative attractions.

Small size of H, N, O and F means atoms on different molecules can come close together before e- clouds repel.

Hydrogen bonding in water explains why water is a liquid at room temperature instead a gas, surface tension (which allows for example bugs to walk on water).

Note empty space in center of ice structure that accounts for why volume of ice is greater than same mass of liquid water and therefore less dense.

Hydrogen Bonding is used to hold together strands of ______.

Characteristics of Hydrogen Bonds that make it useful for holding DNA strands together.

1) Bond Strength – strong enough to hold strands together but weak enough to separate strands without too much energy during replication or gene transcription 2) Specificity – only correct pairing of A with T and C with G positions partial + across from partial to form base pair.

London Dispersion Forces (LDF) – also called Van der Waals:

______dipole attraction between ______molecules.

Momentary shift in electrons in one nonpolar molecule can induce (create) a similar shift in nearby molecules producing temporary dipoles.

Unit 10 ~ Molecular Shapes, Dipoles and Intermolecular Forces~ Assignment List

10-1 Sections: 12.8, 12.9 (p. 363 – 370)

10-2 Sections 12.3 (p. 346, 347) and Dipole Moments Problem Set 12.10 (p. 370- 373)

10-3 Section 14.3 (p. 433 – 435) Page 446, #20-22, 24; Page 447, #26

10-4 Intermolecular Forces Problem Set

10-5 Unit 10 Rev Prob Set (VSEPR/ Dipoles/ IMFs)

10-6 Unit 10 Practice Test

HW 10-1: 1) Why is the three-dimensional shape of a molecule important, especially for biological molecules?

2) How many bonding and nonbonding pairs are present around the CENTRAL ATOM in each of the following molecules?

BP: BP: BP: BP: BP: BP:

NP: NP: NP: NP: NP: NP: - Only count nonbonding pairs on CENTRAL atom. -Double and triple bonds count as only ONE bonding pair (one bonding “region”).

3) The geometry of a diatomic molecule (2 atoms bonded together) is always ______.

4) Do valence electron pairs around the central atom attract or repel? Is the lowest energy arrangement to have the pairs as close together as possible or as far apart as possible?

5) Although both thee BF3 and NF3 molecules contain the same number of atoms, the BF3 molecule is flat whereas the

NF3 molecule is trigonal pyramid. Why are B-F bonds flat in the same plane, whereas N-F are pushed down? (What is pushing the N-F bonds down?

Dipole Moments Problem Set ~ HWK 10-2

Directions: For each of the molecules below: 1) Draw the Lewis Structure. 2) Determine the molecular geometry. 3) Indicate whether or not the molecule has a NET dipole moment. If the molecule does have a net dipole, indicate the direction of the dipole. Electronegativity values: Cl = 3.0, H = 2.1, Be = 1.5, Br =2.8, C = 2.55, S = 2.58, O = 3.5, N= 3.0, F=4.

1) Cl2 2) HCl

3) BeCl2 4) BeClBr

5) CS2 6) H2S

7) NH3 8) BF3

9) CCl4 10) CHCl3

HW 10-3: Intermolecular Forces Problem Set Part 1 (Questions #2,4,6,7 from Zumdahlp. 446, #20-22,24)1)) 1) How does the strength of a typical covalent bond compare to the strength of a typical hydrogen bond, a typical dipole- dipole attraction and a typical London Dispersion Interaction (LDF)?

2) How is the strength of dipole-dipole interactions related to the distance between polar molecules? Are dipole-dipole attractions short-range or long-range forces?

3) For each of the following pairs of molecules draw the correctly oriented dipole-dipole interaction:

A) ICl and ICl B) NH3 and CHCl3 (draw in 3-D, see dipole problem set)

4) What atoms in addition to H are necessary for hydrogen bonding? How does the small size of the hydrogen atom contribute to the unusual strength of the dipole-dipole forces involved in hydrogen bonding?

5) Hydrogen Bonding (Recall H-FON and FON). For each of the following pairs of molecules, draw the lewis dot structure and look for HFON and FON to decide whether or not a hydrogen bond will form. If an H-bond does form, draw the molecules in their correct orientation (can use 2-D).

A) H2O and CH4 H-bond formed?

B) NH3 and NH3 H-bond formed C) HF and CH3OH H-bond formed?

6) Explain the following observation. The boiling point of water, H2O is abnormally high its mass compared to H2S, H2Se and H2Te.

7) Connection to Biology: Hydrogen bonding is often used in biological molecules for example in holding together the strands of DNA molecules. In addition to being a weak enough type of bond to allow the double helix to be unzipped for replication (copying DNA) or transcription (reading the gene as the first step towards making a protein), hydrogen bonds are very specific. In DNA hydrogen help ensure that correct base pair matches are made. In one of the diagrams below, the bases are correctly matched in the other they are mismatched. Which pairs are a correct match and which pair is a mismatch? How do you know? For the correct matches circle the hydrogen bonds.

8) What are London Dispersion forces and how do they arise?

HW 10-4 Relative Strength of IMF’s and Phase Transitions

1A) Rank the 3 types of intermolecular forces [dipole-dipole, hydrogen bond, London Dispersion Force (LDF)], in order according to typical strength relative strength strongest to weakest:

______> ______> ______

1B) What factor can significantly change the relative strength of London Dispersion attractions?

1C) Complete the following: The greater the strength of the intermolecular forces between molecules the (higher/lower) ______the boiling point.

2) For each of the following molecules, identify the strongest class of intermolecular force acting between the atoms or molecules.

A) H2 and H2O B) NH3 and H2O

C)BH2F and BH2F D) BF3 and BF3

E) NH3 and CH3OH F) He and CH4

3) Use your knowledge of intermolecular forces to explain the following observations (Hint: identify the strongest type of intermolecular forces present between molecules).

A) I2 has a higher boiling point than Br2.

Strongest type of IMF between I2 molecules:

Strongest type of IMF between Br2 molecules:

Explanation:

B) H2O has a boiling point than H2S.

Strongest type of IMF between H2O molecules:

Strongest type of IMF between H2S molecules:

Explanation:

4) Rank the following compounds: H2, HF and HCl in order from predicted highest to lowest boiling points: Briefly explain your reasoning in terms in terms of relative IMF strengths.

Strongest type of IMF between H2 molecules:

Strongest type of IMF between HF molecules:

Strongest type of IMF between HCl molecules:

______> ______> ______

Homework 10-5 ~ Unit 10 Review

VSEPR 1) Self-check 12.5, p. 369. Determine geometry name and draw in 3-D.

+ a) NH4 b) NF3 c) H2S

-1 d) ClO3 e) BeF2

2) For each of the molecules below: i) Draw in 3-D ii) Give Geometry Name iii) State the approximate bond angle?

A) CH4 B) NH3 C) H2O

Dipoles 3) What does it mean to say that a molecule is polar or has a net dipole moment?

4) For each of the following molecules, draw the molecules in 3-D and indicate if the molecules will be polar or nonpolar. If a molecule is polar indicate the direction of the net dipole or label the δ- and δ+ sides.

A) AsH3 B) H2S

C) SiH4 D) CS2

Intermolecular Forces 5) Concept of intra vs inter molecular forces. Explain the difference in the boiling points of the following compounds using relative strengths of intermolecular forces for HF and HCl versus the ionic bonding (intramolecular) present in the crystal lattice of LiCl.

HF 20 oC HCl -85 oC LiCl 1360 oC

6) For each substance, identify the strongest type of IMF present and then predict which would have the highest boiling point: HCl, Ar, F2

7) What type of intermolecular force is present in nonpolar molecules? What factor affects the strength of this particular IMF?

8) Explain the following observations: Methane (CH4) and ammonia (NH3) differ in molar mass by only one unit, so their London dispersion force intermolecular forces should be very similar (due to similar size of electron clouds). However, the boiling point of ammonia is over 100 oC higher than that of methane.

9) Explain relative bp values for Noble Gas Family in terms of IMF’s. (From Zumdahl p. 447, #27) He: -272 oC Ne: -245.9 oC Ar: -185.7 oC Kr: : -152.3 oC Xe: -107.1 oC Rn: -61.8 oC

o 10) Explain, in terms of IMF’s, why the mp of ice (frozen H2O) = 0 C, while dry ice (frozen CO2) melts at much lower Temperature. (From Zumdahl p. 448, #46)

Homework 10- 6 Practice Test VSEPR Theory 1) According to the VSEPR model, what determines the geometry of a molecule?

2) What is the geometry name and bond angles for CH4?

3) Place the following in order from largest to smallest bond angles: H2Se, SiCl4, AsH3

4) Draw the following molecules in 3-D using the conventions: For bonds in plane Bonds coming out of plane For bonds going into plane Give bonds angles for A-D.

A) Triangular(trigonal) planar B) Tetrahedral

C) Trigonal bipyramid D) Octahedral

E) Trigonal pyramid

Polar/Nonpolar Molecules 1) What is a polar molecule?

2) Using a diagram and words, explain why a stream of water is attracted to both a + and – charged wand, while a stream of CS2 is not attracted to either wand.

3) Draw a diagram to show how water molecules dissolve a lattice of sodium and chloride ions.

Intermolecular Forces 4) What is an intermolecular force? What is the difference between an intermolecular force and intramolecular force? Which forces are stronger?

5) List 3 examples of intramolecular forces and 3 examples of intermolecular forces.

6) For each of the following molecules, CH2O and CHCl3, draw: A) Lewis Dot Structure B) Correct 3-D structure C) Net dipole D) Correctly oriented dipole-dipole attraction

7A) In the following picture, what type of bond is represented by the solid black line, the dotted line and the dashed line.

7B) If the solid black line represents 100 units of bond energy, what are the approximate values of the dotted and dashed lines?

8) Choose the correct word: The (weaker/stronger) ______the intermolecular forces between two molecules the higher the boiling point.

9) Explain in terms of the intermolecular forces involved the following observations: A) NH3 has a higher boiling point than PCl3.

B) Ne has a lower boiling point than Xe.

C) CHCl3 has a higher boiling point than CCl4.

D) F2 has a lower boiling point than Cl2.

Unit 11 ~ Thermochemistry and Thermodynamics ~ Test Topics Thermochemistry

• Know First Law of Thermodynamics, also called the Law of Conservation of Energy. Be able to apply the concept to the conversion between potential ↔ kinetic energy and in heat transfer problems: -Q = +Q or -[sm∆T] = + [sm∆T] • Understand the concepts of potential energy (stored energy due to position or chemical composition) and kinetic energy (energy of motion). Be able to apply these concepts in explaining experimental observations or demonstrations (like our Ethanol Cannon). • Know the Second Law of Thermodynamics (heat flows from hot to cold) • Understand the concept of specific heat capacity. Specific heat capacity is defined as the number of joules required to raise the temperature of 1 gram of a substance by 1oC. Different substances absorb heat differently. Metals, for example, tend to have low specific heat capacities. This means that metals need to absorb only a small amount of heat in order to increase their temperature. • Know the equation Q = sm∆T and be able to apply this equation to solve heat transfer problems. S values will provided; you do not have to memorize value of s. • Understand the terms system and surroundings as used to describe heat transfer. The system represents the materials undergoing chemical or physical changes, for example, the chemicals reacting in a beaker. The surroundings represent everything else in the universe beside the system. In our example, the beaker would be a part of the surroundings. • Know the terms exothermic and endothermic. Understand that exothermic processes release heat energy and that the sign of q or ∆H is -. Understand that endothermic processes absorb heat energy and that the sign of q or ∆H is +. The sign is always from the prospective of the system. • Be able to draw endothermic and exothermic reaction profile diagrams, indicating the relative positions of Reactants and Products. Recognize that the higher the energy, the less stable and weaker are the bonds, whereas low energy represents more stable, stronger bonds. • An endothermic process would cause the outside of the beaker to feel cold. • An exothermic process would cause the outside of the beaker to feel hot. • Understand that breaking bonds requires energy and making bonds releases energy. The overall enthalpy change for a reaction indicates whether the bonds of the product are lower in energy (more stable) than the bonds in the products or vice versa. • Be able to calculate the enthalpy change for a reaction that occurs when a certain number of grams or moles of chemical react or are produced. Balanced equation and bond energy values will be provided.

Heat Transfer Problem

100.0 g of water at a temperature of 22.4 oC is placed in a calorimeter. A 75.25 g sample of Al is removed from boiling water at a temperature of 99.3 oC and quickly placed in the calorimeter. The substances (Al and water) reach a final temperature of 32.9 oC. The specific heat capacity of water is 4.18 J/g oC. Determine the specific heat capacity of the aluminum. Experimental setup:

o Boiling water bath at 99.3 C Aluminum cylinder, 75.25 g 100.0 g of water at 22.4 oC What is a calorimeter?

Steps for solving problem: Translate information from the problem into a data table

Variable/Units Aluminum Water (in calorimeter) s (J/goC) m (mass in grams)

o Tfinal (Final Temperature in C)

o Tinitial (Initial Temperature in C)

∆T (Tf – Ti)

Trace Heat Flow (Using common sense and laws of thermodynamics): 2nd Law of Thermodynamics: Heat always flow from hot to cold: in our example this is from ______(99.3oC) to ______in calorimeter (22.4 oC). 1st Law of Thermodynamicss: Heat energy is conserved: Heat Lost by ______= Heat Gained by ______(Assume no heat lost to surroundings) (Recall Symbol for heat is Q, thus in equation form:

Since Q = sm∆T, substituting into the above equation:

-[ sm∆T] Al = +[ sm∆T] H2O Substituting in from table on previous page:

Variable/Units Aluminum Water (in calorimeter) s (J/goC) ? 4.18 m (mass in grams) 75.25 g 100.0

o Tfinal (Final Temperature in C) 32.9 32.9

o Tinitial (Initial Temperature in C) 99.3 22.4

∆T (Tf – Ti) -66.4 10.5

o o o o o -[ sAl( g)( C - C)] = +( J/g C)( g)( C – C) o - S Al (75.25)(-66.4) g C = (418)(10.5) J Notice the substance that loses heat (-Q) will a - ∆T. Two negatives make a positive! o S Al (4996.6) g C = 4389 J o S Al = 4389 J = 0.87787 → 0.878 J/ g C 3 SF o 4996.6 g C

Notes ~ Enthalpy, Endothermic and Exothermic Reactions

Enthalpy – a measure of the flow of ______under conditions of constant pressure. Describing the direction of heat energy flow:

System: The part of the ______; in a chemical reaction the system is the ______.

Surroundings: ______of the ______which is not the system; in our demonstrations of chemical reaction, the surroundings would include the beaker, the desk top, the air in the room, etc.

Exothermic process- R → P + ______∆H = heat energy ______(exits) the system; Heat flows from the ______to ______. Exothermic rxns feel ______: Temp rxn Temp surroundings

Endothermic process- R + ______→ P ; ∆H = heat energy ______(enters) ______from ______. Endothermic rxns feel ______Temp rxn Temp surroundings

Key idea: 2nd law of Thermodynamics: Heat always flows from ______to ______. Exothermic and Endothermic are always defined from the ______of the ______. What do the chemicals need to do (______or ______) to equilibrate with the surroundings and return to room temperature?

Diagram Exothermic Reaction Diagram Endothermic Reaction

For each demonstration: Label the system and the surroundings, describe whether the reaction feels hot or cold, compare the temperature of the system to the surroundings, indicate the direction of heat flow (system → surroundings or surroundings → system) and label as exothermic or endothermic.

Demonstration Example #1: NH4Cl in water:

Chemicals (NH4Cl + H2O) = Beaker, desk top, air, you = Observations: Rxn feels . T chemicals T surroundings; Heat flows from to ; Reaction is ∆H =

Demonstration Example #2: CaCl2 in water:

Chemicals (CaCl2+ H2O) = Beaker, desk top, air, you = Observations : Rxn feels . T chemicals T surroundings; Heat flows from to ; Reaction is ∆H =

Concept Check: Which graph of temperature vs time represents an exothermic reaction and which represents an endothermic reaction?

Temperature Temperature

Room Temp Room Temp

Time Time

Homework 11-1 Energy Concepts Problem Set Energy = capacity to do ; Work = Working” defn of energy = Potential Energy – Examples: Positional: Chemical composition: Hi PE = bonds, unstable; Low PE = Kinetic Energy – First Law of Thermodynamics –

1) Thermochemistry is the study of heat energy changes in chemical reactions. State two reasons it might be important to be able to predict the heat energy changes associated with a particular chemical reaction and provide supporting examples.

2) Imagine a ball is held in the air 5 ft above the ground and then released. Use the First Law of Thermodynamics to complete the table below.

Location PE KE Total Position #1 (Release point) Units Units Energy Units Position #1 10 0 10 Position #2 (ball falling)

Position #2 6 10 Position #3 (ball falling) Position #3 8 Ground

3) The “ethanol cannon” demonstration also illustrates the transformation of potential energy into kinetic energy. A small quantity of ethanol (CH3CH2OH) is placed in a sealed shampoo bottle fitted with a rubber stopper and allowed to vaporize (change from liquid to gas). A) Potential is energy is stored energy due to position or chemical composition. Why does the ethanol and oxygen mixture have potential energy?

B) What type of reaction pattern is this? What are the products of the reaction?

Reaction Type: CH3CH2OH + O2 → C) This reaction is thermodynamically favorable, (meaning it should take place), however no reaction is initially observed. Explain why no reaction is observed and what must be provided to allow the reaction to occur.

PE D) Using to represent units of chemical potential energy and K E to represent units of

Kinetic energy of molecules, fill in the missing number of K E for the products. (Hint: Remember The Law of Conservation of Energy) REACTANTS: PRODUCTS:

PE PE PE PE KE KE PE

KE E) Complete the following table comparing the relative reactants and products by filling in the correct description from the word bank below: High Potential Energy (Hi PE) , Low Potential Energy (Low PE), High Kinetic energy (Hi KE), Low Kinetic Energy (Lo KE), Stable, Unstable. Stronger chemical bonds, weaker chemical bonds.

Property Reactants Products Potential Energy Kinetic Energy

Stable/Unstable

Stronger bonds/Weaker bonds

F) Sketch a reaction profile diagram for the reaction. Label Reactants, Products, and transition state (TS).

Potential energy

Reaction Progress 98

HW 11-2 and 11-3 Temperature and Heat Key Ideas 2nd Law of Thermodynamics

Temperature-

Heat-

Homework 11-2 – Concept of Temperature Imagine a container with holding 5 Helium atoms. Use the data table below to answer questions #1A, 1B

Molecule # Kinetic Energy ( J) 1 15 2 5 3 10 4 8 5 12

1A) Do all of the atoms have the exact same kinetic energy?

Draw of bar graph showing Kinetic Energy on the Y axis and molecule # on the X-axis.

Kinetic Energy (J)

Molecule #

1B) Calculate the average kinetic energy for the set of molecules.

1C) Concept of Temperature: The temperature of a set of molecules is directly proportional to the average ______of the molecules. 99

2) Applying the Second Law of Thermodynamics: In the diagram below each box represents a molecule and the value inside the box represents the kinetic energy of the molecule. System #1 initial System #2 initial

KE = 5 KE = 4 KE = 2 KE = 6 KE = 15 KE = 10 KE = 4

2A) Calculate the average KE (units are J) for each system: Average KE (system #1): Average KE (system #2): 2B) Which system is at the higher temperature? 2C) The 2nd Law of Thermodynamics states that heat allows flows from _____ to ______. 2D) If system #1 and system #2 are allowed to come in direct physical contact which system will gain heat energy? Which will lose heat energy?

2E) The diagram below represents the two systems after THEY HAVE BEEN IN CONTACT WITH EACH OTHER FOR AN EXTENDED PERIOD OF TIME. System #1 System #2

KE = 6 KE = 7 KE = 6 KE = 7 KE = 8 KE = 8

2F) Calculate the average KE (units are J) for each system: Average KE (system #1): Average KE (system #2): 2G) Compare the temperatures (average KE) values for each system. Why do we describe this condition as thermal equilibrium?

2H) Predict the general shape of a temperature vs time graph that would result system when system #1 and system #2 have been placed in contact with each other. The first pair of data point temperatures at time zero, when the systems are first contact each other, is provided. Label the onset of thermal equilibrium.

Key: Temperature Temp. system 1 molecules =

Temp. system 2 molecules =

Time

Initial contact 100

HW 11-3 Concept of Heat: Developing and Understanding of the Heat Energy Equation and relationship between Heat and Temperature Heat is a transfer of energy that occurs from a hotter object to a cold object. While heat energy transfer is directly linked to temperature change, ΔT, it is NOT identical to temperature. The next several problems will explore the relationship between temperature and heat. 1) If an equal amount of heat energy is transferred to a cup of water and a billion gallons of water (e.g. Niagara Falls), which of the following statements would be true? A) The final temperatures of the cup of water and a billion gallons of water would be exactly the same, because each absorbed the exact same amount of heat energy. B) The cup of water would be at a higher temperature than the billion gallons of water. C) The billion gallons of water would be at a higher temperature than the cup.

2A) If $2000 were divided equally between 2000 students, how much money $ would each student receive?

If the same $2000 were divided equally between 20 students, how much $ would each student receive?

2B) Imagine that equal amounts of heat energy (10 energy units), were transferred to a container with two water molecules and a separate container with five water molecules. The water molecules in each cup were initially as the exact same temperature. Using the symbol E to represent an energy unit, draw a diagram to show how many energy units each water molecule would absorb.

Cup #1 Cup #2 Are the cups at the same temperature or different temperatures? If different, would cup #1 or cup #2 be at the higher temperature? Explain in terms of average kinetic energy per molecule.

3) Imagine two containers. The first holds five $100 dollar bills. The second holds one hundred $50 dollar bills.

Container #1: Container #2: $50 X 100 $100 $100 $100 $100 $100

3a) In which container are individual bills worth more?

3b) What is the total amount of money in each container? 101

4) Imagine a bucket of water at 50 oC and a cup of boiling water 100 oC. 4A) In which container, the bucket or the cup, do the water molecules have the highest average kinetic energy?

4B) Which container, the bucket of water at 50oC or the cup of boiling water at 100 oC, could transfer the greatest amount of heat energy to a cold water bath? Explain.

5)The kinetic energies of molecules present two different containers are given below. Container #1 Container #2

KE = 6 KE = 4 KE = 2 KE = 6 KE = 4 KE = 2 KE = 4

KE = 4 KE = 5 KE = 3

5A) Calculate the AVERAGE kinetic energy of the molecules in each container: Container #1: Container #2 5B) Calculate the TOTAL kinetic energy of the molecules in each system: Container #1: Container #2:

5C) Are the temperatures of the two containers the same or different? If different, which container has a higher temperature? Explain.

5D) Is the total amount of heat energy stored in the two containers the same or different? If different, which container has more stored heat energy? Explain.

102

6)The kinetic energies of molecules present two different containers are given below. Use the given information to answer questions. Container #1 Container #2

KE = 120 KE = 100 KE = 50 KE = 48 KE = 52

KE = 80 KE = 45 KE = 50 KE = 55

KE = 40 KE = 60 KE = 50

KE = 50

6A) Calculate the AVERAGE kinetic energy of the molecules in each container: Container #1: Container #2: 6B) Calculate the TOTAL kinetic energy of the molecules in each system: Container #1: Container #2:

6C) Are the temperatures of the two containers the same or different? If different, which container has a higher temperature? Explain.

6D) Is the total amount of heat energy stored in the two containers the same or different? If different, which container has more stored heat energy? Explain.

103

Concept of Specific Heat Capacity: Do all objects absorb and store heat energy in exactly the same way? 7) Defn Specific Heat Capacity: s =

8) Explain the following observations using the concept of specific heat capacity. The temperature of a rock at 3 pm the afternoon on a summer’s day is 120 oF; The temperature of the same rock at 5 am the next day is 65 oF. The temperature of ocean water off the coast of Florida at 3 pm the afternoon on a summer’s day is 95 oF; The temperature of the same water at 5 am the next day is 89 oF.

Given: Rock, s = 0.60 J/g oC Water, s = 4.18 J/g oC

9) What is the equation for heat energy Q?

10) Calculate the heat energy, Q, absorbed when 454 g of water is heated from an initial temperature of 5.4 oC to 98.6 oC. Given: The specific heat capacity of water is 4.18 J/g oC.

104

Homework 11-4 Exploring Concept of Heat vs. Temperature with Silver and Steel Spoons Case #1: If the temperature of 2 objects changes by the same amount, does this mean each object has gained or lost the exact same amount of heat energy? Two 10.0 g spoons, one made of silver, and one made of steel initially at a temperature of 25 oC, are placed into beakers of water maintained at 100. oC with a hotplate. The specific heat capacity of silver, s = 0.23 J/ g oC. The specific heat capacity of steel, s = 0.51 J/g oC. 1) Draw a graph with temperature on the y-axis and time on the x-axis. Draw a heating curve for each metal on the same graph.

Temperature, oC

Time (min) 20 min

How are the curves for each metal alike, how are they different and why?

Use these choices to answer questions #2-4: A) Silver B) Steel C) Both are equal

2) After 20 minutes, which spoon will be at a higher temperature? Explain.

3) After 20 minutes, which spoon’s atoms will have a greater average kinetic energy? Explain.

4) After 20 minutes, which spoon will have stored more heat (Q)? Support your answer with a calculation of o heat absorbed for each spoon. (Hint: Ti = 25 C for each; what is Tf after 20 minutes? Equation: Q = smΔT)

Silver: Steel: o o o o s = 0.23 J/g C m = 10.0 g Ti = 25 C Tf= s = 0.51 J/g C m = 10.0 g Ti = 25 C Tf=

Q= Q=

105

Case #2 - If two objects absorb the exact same amount of heat energy in joules, will they be at the same temperature? The same 10.0 g silver and stainless steel spoons are both at room temperature (25oC). Each spoon absorbs 1000. J of heat energy. Calculate each spoon’s final temperature. (Hint: Solve Q = sm∆T for ∆T, then use ∆T = Tf – Ti) Silver: Steel

9) 1 g samples of copper, Cu and aluminum, Al are allowed to sit in identical 100 oC boiling water baths until their temperatures have equilibrated with the boiling water bath (i.e. both metals are at exactly 100.0 oC). The metals are then transferred to two separate perfect calorimeters (assume no heat escapes from container – all heat from metal is absorbed the water) that both contain the exact same mass of water (10.0 g) at the exact same initial temperature (25.0 oC). Given the following specific heat capacities: Copper, Cu: 0.39 J/g oC Aluminum, Al: 0.89 J/ g oC Sketch a temperature vs. time graph for the WATER in the calorimeter containing each metal on the same graph below. Assume that the calorimeter is perfect (no heat ever escapes out of the calorimeter.) Clearly label each graph (Cu or Al).

Homework 12-5 Heat Transfer Problem Set

106

HW 11-5 Heat Transfer Problems Given: Specific Heat capacity of water = 4.18 J/g oC; Density of water = 1.00 g/mL 1) 50.0 mL of water at a temperature of 24.0 oC is placed in a calorimeter. A piece of unknown metal with a mass of 23.8 g is heated to 100.0oC into the water in the calorimeter. The final temperature of the system equilibrates to 32.5 oC. What is the specific heat of the metal?

Variable Metal Water s (J/ g oC) 4.18 m (g) o Tf C o Ti C ΔT

Since Ti metal > Ti water : - Q metal = + Q H2O

2) A geologist at a mining company is trying to identify a metal sample obtained from a core sample. The unknown metal, with a mass of 5.05 g is heated to 100.00 oC and dropped into 10.0 mL of water at 22.00 oC. The final temperature of the system is 23.83 oC. What is the specific heat capacity of the metal? Variable Metal Water

3) A blacksmith heated an iron bar (s = 0.449 J/ g oC) to 1445 oC. The blacksmith then tempered the metal by dropping it into 4.28 x 103 mL of water that had a temperature of 22 oC. The final temperature of the system was 45 oC. What is the mass in grams of the iron bar?

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4) A 752 mL sample of water was placed in a 1.00 kg (Recall 1.00 kg = 1.00 x 103 g) aluminum pan (s = 0.89 J/ g oC). The initial temperature of the pan was 26.0 oC and the final temperature of the system was 39.0 oC. Which statement could be true about the initial temperature of the water? A) The initial temperature of the water could have been 35 oC. B) The initial temperature of the water could have been 39 oC. C) The initial temperature of the water could have been 43 oC. Explain:

5) 100. g of iron (s = 0.449 J/oC) at a temperature of 100 oC is dropped into 100.0 g of water (s = 4.18 J/g oC) in a perfect calorimeter. The initial temperature of the water in the calorimeter is 20.0 oC. Which of the following statements will be true when thermal equilibrium is reached? A) The thermal equilibrium temperature will be greater than 60 oC. B) The thermal equilibrium temperature will exactly equal to 60 oC. C) The thermal equilibrium temperature will be less than 60 oC. Explain.

6) 100. g of water (s = 4.18 J/oC) at a temperature of 100 oC is dropped into 100.0 g of water in a perfect calorimeter. The initial temperature of the water in the calorimeter is 20.0 oC . Which of the following statements will be true when thermal equilibrium is reached? A) The thermal equilibrium temperature will be greater than 60 oC. B) The thermal equilibrium temperature will exactly equal to 60 oC. C) The thermal equilibrium temperature will be less than 60 oC. Explain.

Answers to Calculations: 1) 1.1 J/ g oC 2) 0.199 J/g oC 3) m = 650 g

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Homework 11 – 6 Endothermic/Exothermic Problem Set 1) Demonstration of potassium permanganate and glycerin reaction:

14KMnO4 + 4C3H5(OH)3 --> 7K2CO3 + 7Mn2O3 + 5CO2 + 16H2O

A) Initial Observations:

Final Observations:

B) At the completion of the reaction, which diagram shows what will happen terms of heat flow?

i) ii)

Does heat have to flow into the system or out of the system to return the system to room temperature?

C) Is this reaction endothermic or exothermic? ______What is the sign of ∆H? ______D) Which best describes this reaction? i) Heat + R → P (heat as reactant) ii) R → P + Heat (heat as product)

E) Sketch a reaction profile diagram for this reaction, labeling the relative positions of R and P.

Reaction Progress

2. Use the demo from the video and the following equation to answer the parts of this question: Ba(OH)2 (s) ∙ 8 H2O(l) + 2 NH4SCN(s) → Ba(SCN)2 (aq) + 2 NH3 (g) + 10 H2O(l) A) Initial observations - Draw a labeled diagram that includes the beaker, chemicals, puddle of water on board

B) final observations: 109

C) At the completion of the reaction, which diagram shows what will happen terms of heat flow?

i) ii)

Does heat have to flow into the system or out of the system to return the system to room temperature?

D) Are the chemicals inside the beaker part of the system or the surroundings?

E) Is the puddle of water of underneath the beaker part of the system or the surroundings?

F) Explain, in terms of heat flow, system and surroundings, why does the beaker stick to the board?

G) Is this reaction endothermic or exothermic? ______What is the sign of ∆H? ______D) Which best describes this reaction: i) Heat + R → P (heat as reactant) ii) R → P + Heat (heat as product) E) Sketch reaction profile diagrams for each Reaction

Reaction Progress 3.) Using the ideas that endothermic (∆H = +) absorbs heat and exothermic (∆H = -) releases heat, indicate whether the reaction is endo or exo and then predict the sign of ∆H:

A) H2O(l) → H2O(g) ∆H = _____ (Hint: This equation represents evaporation. Do you feel hot or cold when water evaporates from your skin?)

B) 2 H2 (g) + O2 (g) → 2 H2O(g) ∆H = _____ (Hint: The Hindenberg Balloon Demonstration from first semester)

C) 2 H2O(g) → 2 H2 (g) + O2(g) ∆H = _____

(Hint: This is our electrolysis demonstration from first semester. Notice it is the reverse reaction from D) 110

Homework 11-7 Breaking and Making Bonds Problem Set

1) BREAKING chemical bonds ______ENERGY;

MAKING chemical bonds ______ENERGY.

2) Most chemical reactions involve both making and breaking chemical bonds. However, in some reactions bonds are only formed or broken, making it easy to predict whether the reaction will be endothermic or exothermic. For each of the following reactions: i) Diagram the reaction using to represent atoms. ii) Determine whether bonds were broken or formed. iii) Predict the sign of ∆H. A) A reaction that occurs in the stratosphere as sunlight hits an oxygen molecule is: O2 (g) → 2 O(g) ∆H = _____

B) Another reaction in the stratosphere is Cl (g) + Cl (g) → Cl2 (g) ∆H = _____

3) In most chemical reactions, bonds are both being broken and formed.

A) The NET ENERGY CHANGE FOR A REACTION is the ______between the energy ______when old bonds are broken and the energy ______when new bonds are formed. B) Consider a hypothetical reaction: A-B + C → A-C + B If bond strength of A-B = 150 units and A-C bond strength = 100 units i) Which is stronger the bonds in the reactant A-B or the product A-C? ii) What is the net energy change for the reaction? (Hint: Breaking costs energy= + sign, making bonds releases energy = – sign)

iii) Is this reaction endothermic or exothermic?

C) In an ENDOTHERMIC REACTION, the reactants have (stronger or weaker) ______bonds and are therefore (more or less) ______thermodynamically stable than the products. The difference in energy between reactants and products is (absorbed or released) ______.

For an ENDOTHERMIC reaction, is the energy term written on the reactant side or the product side? ______+ reactants → products + ______

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D) Consider a hypothetical reaction: X-Y + Z → X-Z + Y If bond strength of X-Y = 50 units and X-Z bond strength = 200 units i) Which is stronger the bonds in the reactant X-Y or the product X-Z? ii) What is the net energy change for the reaction? (Hint: Breaking costs energy= + sign, making bonds releases energy = – sign)

iii) Is this reaction endothermic or exothermic?

E) In an EXOTHERMIC REACTION, the reactants have (stronger or weaker) ______bonds and are therefore (more or less) ______thermodynamically stable than the products. The difference in energy between reactants and products is (absorbed or released) ______.

For an EXOTHERMIC reaction, is the energy term written on the reactant side or the product side? ______+ reactants → products + ______

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Enthalpy (∆H) and Stoichiometry Problem Set ~ Homework 11-8 Notes: Balanced equation links moles of reactants, products, and heat energy absorbed or released

2 H2 + O2 → 2 H2O ∆Hrxn = -572 kJ (exothermic) can also be written as 2 H2 + O2 → 2 H2O + 572 kJ means 2 moles H2 + 1 mole O2 → 2 moles H2O + 572 kJ

1 mole H2 + 1/2 mole O2 → ____ moles + ______kJ

4 moles H2 + 2 moles O2 → ____ moles + ______kJ

Example Problem: Calculate the amount of heat released when 8.08 g H2 burn in excess O2.

Directions for #1 – 3: • Using the ideas of stoichiometry (coefficients represent mole ratios) and dimensional analysis that we learned last semester, solve for the heat. Since answers are provided, you MUST SHOW WORK for credit!

1) Given the thermite reaction: 2 Al + Fe2O3 → Al2O3 + 2 Fe ∆H = -850 kJ Calculate the amount of heat energy released by the reaction of 9.50 g of Al.

2) Given the reaction: 4 NH3 + 7 O2 → 4 NO2 + 6 H2O ∆H = -1396 kJ Calculate the amount of heat energy released by: a) The reaction of 48.0 g of oxygen.

b) The formation of 75.0 g of NO2.

3) Given the reaction: 2 SO3 → 2 S + 3 O2 ∆H = +395.2 kJ Calculate the amount of heat energy absorbed by the reaction of 50.0 g of SO3.

Answers: 1) -150. kJ 2) -299 kJ 3) -569 kJ 4) +123 kJ

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Entropy Concept – Homework 11-9 Entropy(S) – is a measure of ______or randomness ∆S = change in entropy ; ∆S = + , entropy increasing ; ∆S = - , entropy decreasing Spontaneous process – a process that occurs without outside intervention, i.e. something that happens “ on its own” 2nd Law of Thermodynamics – the entropy of the universe always ______for a spontaneous process Entropy – is determined by the ______of possible arrangements or microstates Local order can be imposed only by doing work, e.g. you can organize your room by doing work. Concept of Spontaneous – which of the following events will naturally occur? (Yes or No) 1) A dinner plates drops and shatters into many irregular pieces.

2) Over a few weeks, your room cleans itself; clothes fold themselves and migrate to drawers.

3) A rock rolls itself up a cliff and balances on a ledge.

4) The random motion of N2 and O2 molecules an entire classroom of air to end up in one corner of the room, leaving the students in a vacuum.

Concept of number of microstates or arrangements- which has more possible ways to arrange things? 1) Socks folded together with their matched partner placed in the top drawer of your dresser vs individual socks hanging from the lamp and desk chair, under your bed, in your bed.

2) Atoms in the solid vs atoms in the liquid phase vs atoms in the gas phase

Predicting the Entropy changes: For each of the following processes predict whether entropy will increase, ∆S = + (more disordered) or decrease, ∆S = -, less disordered.

3) H2O(l) → H2O(s)

4) CO2 (s) → CO2 (g)

5) We saw a laser disc segment that has Don Showalter mixing two white solids in a beaker that freezes to a board. Use the phase of matter of the reaction (as seen below) to predict the entropy change.

Ba(OH)2∙8 H2O(s) + 2 NH4SCN(s) → Ba(SCN)2(aq) + 2 NH3(aq) + 10 H2O(l)

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HW 11-10 ~ Thermochemistry/ Thermodynamics ~Review Problem Set Concept of 1st Law of Thermodynamics – Conversion of Chemical Potential Energy into Kinetic Energy

1) Gasoline, consists mainly of octanes, C8H18. Gasoline can react with oxygen O2 in the following reaction.

2 C8H18 + 25 O2 → 16 CO2 + 18 H2O A) This reaction is thermodynamically favorable, meaning that the reaction should spontaneously occur. When gasoline is in the gas tank of your car and is in contact with air inside the engine, why does no reaction take place? What must you do to start the engine?

B) True or False ~ Which of the following statements about this reaction are true? I) The gasoline has potential energy due to its chemical composition.

II) The products of the reaction, CO2 and H2O, have higher potential energy but lower kinetic energy than the reactants, C8H18 + O2.

III) The reactants, C8H18 + O2, have weaker, less stable bonds than the products, CO2 and H2O. C) Sketch a reaction profile diagram for this reaction, with Potential Energy on the y-axis and Reaction Progress on the x- axis. Label reactants, products and transition state. Is this reaction endothermic or exothermic?

Concept of Heat vs. Temperature 2A) Imagine a drop of boiling water falls on your skin. What is the temperature of this water (at sea level)?

2B) Now imagine you knock over a cup of boiling water on your skin. i) How does the temperature of boiling water in the cup compare to that of the drop?

ii) Why does the cup of water hurt soooo much more than the drop?

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3) Calculate the quantity of heat that must be removed (released) from 200.0 g of ethyl alcohol to cool it from 25.0 oC to 10.0 oC . ( s for ethyl alcohol = 2.45 J/g oC)) [ Ans: -7350 J]

4) Consider two teaspoons, each having a mass of 10.0 g. One is made of nickel (specific heat of 0.440 J/g◦C) and the other of plastic (specific heat of 1.30 J/g◦C). They are immersed in boiling water at 100.0 oC.

A) After 20 minutes in the boiling water, which spoon will have a higher temperature? Explain.

B) After 20 minutes in the boiling water which spoon will stored have a greater amount of heat? Explain.

C) Calculate how much heat would be released if the two spoons were each cooled separately to 25.0 ◦C (assume the boiling water is at 100.0 ◦C). Be sure to include the correct sign of Q, SHOW WORK, and round final answer to 3 Significant Figures.

Nickel spoon: Plastic spoon: s = 0.440 J/g oC s = 1.30 J/g oC m = 10.0 g m = 10.0 g

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5) The same two teaspoons from the previous question, each with a mass of 10.0 g, each absorb the exact same amount of heat energy, 75.0 J. A. If both spoons start at the same initial temperature, 20.0 ◦C, which statement will be true? I) The nickel spoon will be at a higher temperature than the plastic. II) The plastic spoon will be at the higher temperature than the nickel. III) Since the have absorbed the exact same amount of heat both spoons will be the same final temperature. B. Calculate the final temperature each spoon. Was your answer to part A correct? Nickel spoon: Plastic spoon: s = 0.440 J/g oC s = 1.30 J/g oC m = 10.0 g m = 10.0 g

6) 1 g samples of gold, Au and zinc, Zn are allowed to sit in identical 100 oC boiling water baths until their temperatures have equilibrated with the boiling water bath (i.e. both metals are at exactly 100.0 oC). The metals are then transferred to two separate perfect calorimeters (assume no heat escapes from container – all heat from metal is absorbed the water) that both contain the exact same mass of water (10.0 g) at the exact same initial temperature (25.0 oC). Given the following specific heat capacities: Gold, Au: 0.13 J/g oC Zinc, Zn: 0.39 J/ g oC After 20 minutes which statement will be true? A) The temperature of the water in each calorimeter will be exactly the same. B) The temperature of the water bath containing the gold will be higher than the bath containing the zinc. C) The temperature of the water bath containing the zinc will be higer than the bath containing the gold. Explain:

7) A thin disk of solid silvery Gallium metal (melting point = 86oC), is placed on a board. A beaker is placed on top of the gallium disk. Purple Potassium permanganate crystals are placed into the beaker and Glycerin, a thick syrupy clear, colorless liquid is poured on top. A cloud smoke is slowly produced, followed by bright white sparks. The beaker feels very hot to the touch. When the beaker is lifted off the board it is observed that the silvery metal on the board is now a puddle of silvery liquid. A) What is the system?

B) Is the reaction exothermic or endothermic? Describe the direction of heat flow in terms of system and surroundings?

C) Why is silvery metal (Gallium) now a liquid? Is the gallium part of the system or the surroundings?

D) Sketch a reaction profile diagram of for this reaction with enthalpy on the y-axis and reaction progress on the x-axis. Clearly Label Reactants and products.

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8) In photosynthesis sunlight is used to convert CO2 and H2O are converted into glucose and oxygen.

6 CO2 + H2O + light → C6H12O6 + 6 CO2 Is this reaction endothermic or exothermic? What is the sign of ∆H? 9) 4 g of pink cobalt(II) chloride is placed in a test tube and the temperature is measured at 21oC. 20 mL of thionyl chloride, a clear colorless liquid is added. A gas is given off and the color changes from pink to blue. After 7 minutes, the temperature of the test tube has dropped to 6 oC. A) After the reaction, to return to room temperature, will heat have to flow into the system or out of the system?

B) What is the sign of ∆H? Is the reaction endothermic or exothermic? C) Sketch a reaction profile diagram with potential energy or enthalpy on the y-axis and reaction progress on the x-axis. Label reactants, products and transition state.

10) Predict the sign of ∆H for each of the following:

A) Br + Br → Br2 ∆H = B) F2 → F + F ∆H = C) Br2 (liquid) → Br2 (gas) ∆H =

11) Fill in the blanks: Entropy is a measure of the amount of ______(order or disorder) in a system. The greater the number of possible positions of atoms or molecules, the ______(greater or lower) the entropy.

12) The reaction between glycerin and potassium permanganate exists in the following phases for reactants and products: Solid + Liquid → solid + solid + liquid + gas

Is the ∆S for this reaction + (entropy increasing) or – (entropy decreasing). How do you know?

13) Liquid hydrogen peroxide (H2O2) is an oxidizing agent in many rocket fuel mixtures because it releases oxygen gas on decomposition: 2 H2O2(l) → 2 H2O(l) + O2(g) ΔHrxn = -196.1 kJ A) Is this reaction endothermic or exothermic? B) If the reaction was rewritten so that the heat term (196.1 kJ) was included in the reaction would it be placed on the reactant side or the product side?

C) If 1 mole of H2O2 reacts, how much moles of H2O and O2 would be produced and how much heat would be produced?

D) How much heat (in kJ) is released when 150 g of hydrogen peroxide decompose?

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Unit 12 pH mini Unit ~Test Topics pH and Acid/Base Solutions - Know definition of pH = -log[H+]; understand for logarithmic scales such as pH, each 1 unit difference in pH represents a factor of 10 difference in [H+] - Understand inverse relationship between [H+] and [OH-]: [H+][OH-] = 1 x 10-14 - Know definition of acidic, neutral, basic and solutions in terms of [H+], [OH-] and pH values. Acidic solutions [H+] > [OH-] and pH < 7 Neutral solutions [H+] = [OH-] and pH =7 Basic solutions [H+] < [OH-] and pH > 7 - Understand importance of pH in impacting types of reactions and rates of reactions that can occur in aqueous solutions; examples unfolding of proteins and effecting concentration of heavy metals ions. pH problem Set 1A) Enzymes are proteins that catalyze all of the chemical reactions in your body. Why can changing the pH of your blood be fatal? Draw a picture of a normal enzyme and a second picture of an enzyme in an acidic or basic solution to support your answer.

1B) Why does changing the pH of water from mines in the Colorado mountains impact the quality of drinking water in towns like Leadville?

2)Acids are H+ ______; Bases are H+ ______

3)Define Acidic, Basic and Neutral Solutions in terms of [H+],[OH-] and pH by filling in the missing information. In the boxes show the relative amounts of H+ vs. OH- ions that would be present in each type of solution by drawing 10 total ions.

Neutral Solutions Acidic Solutions Basic Solution [H+ ] [OH-] [H+ ] [OH-] [H+ ] [OH-] pH 7 pH 7 pH 7

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4A)Fill in the missing information: Note [ ] represents Molar concentrations (moles/ liter) + - Solution [ H ] [OH ] Letter A -1 1 x 10 1 x 10-13 B 1 x 10-2 1 x 10-12 C 1 x 10-3

D 1 x 10-10 E 1 x 10-7 1 x 10-7 F 1 x 10-9 G 1 x 10-10 H 1 x 10-14

4B) What is the relationship between [H+ ] and [OH-]?

i) Circle correct choice: As [H+] ↑, [OH-] ↑ or ↓

ii) Mathematical relationship: [H+][OH-] =

5A) What is the variation (range) in the concentration of H+ typically found in aqueous (water) solutions?

5B) Why do chemists use a logarithmic scale to describe the pH of a solution instead stating the exact molar concentration of H+ ion?

5C) What is the definition of pH?

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6A) Fill in the missing information: 10X 10X 10X 10X 10X 10X pH = 1 2 3 4 5 6 7 [H+] = 0.1 0.01 0.001 [H+] = 1 x 10-1 1 x 10-2 1 x 10-3 1 x 10- 1 x 10- 1 x 10- 1 x 10-

6B) In the logarithmic pH scale. each unit difference in pH corresponds to a factor of _____ in [H+]. 6C) Fill in the missing information / circle the correct choice: i) Compared to a solution with a pH of 2, a solution with a pH of 1 has [H+] which is ______X (higher or lower).

ii) Compared to a solution with a pH of 3, a solution with a pH of 1 has [H+] which is ______X (higher or lower).

iii) Compared to a solution with a pH of 4, a solution with a pH of 1 has [H+] which is ______X (higher or lower).

7A) Which has a higher [H+], a solution with a pH of 7 or a solution with a pH of 9?

7B) By what factor [H+], of are the concentrations different?

8A) Which has a higher [H+], a solution with a pH of 8 or a solution with a pH of 5?

8B) By what factor [H+], of are the concentrations different?

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9A) Fill in the missing information:

Solution [H+] pH Letter A 1 x 10-1 1 B 1 x 10-2 2 C 1 x 10-3 3 D 1 x 10-4 E 1 x 10-7 F 1 x 10-9 G 1 x 10-10 H 1 x 10-14 14

9B) List the letters of the above solutions which are acidic.

9C) List the letters of the above solutions are basic.

9D) List the letter of the above solutions which is neutral.

9E) Which of the above solutions have [H+] > [OH-]?

9F) Which of the above have [OH-] > [H+]?

9G) Which of the above have [H+] = [OH-]?

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Required Practice Problems for 2nd Semester Final Exam Use your own paper Unit 7 Stoichiometry Given the following reaction:

2 Al + 3S → Al2S3

3A) Calculate the theoretical yield of Al2S3 that could be formed from the reaction of 10.0 g of Al with 15.0 g of S. Which is the limiting reactant?

3B) If 21.7 g of Al2S3 is actually obtained from the reaction in the lab, calculate the % yield. Unit 8 – Quantum and the Periodic Table 4) List 3 postulates of the Bohr model and state which one is NOT part of modern quantum theory. 5) What does the Heisenberg Uncertainty Principle state regarding the precise path of an electron? 6) What is an orbital? 7) Write the complete electron configuration and orbital notation diagram for Mg, element #12. 8) Write the abbreviated (core notation) electron configuration and orbital notation for the valence electrons only for Ga, element #31. 9) Define valence electrons. How do you determine the number of valence electrons for an element. What is true about the number of valence electrons in a group? What does that imply for the chemical properties of elements within a group? 10) How does atomic radius vary across a period? Down a group? How does ionization energy vary across a period? Down a group?

Unit 9 Bonding through Lewis Structures 11) What is an ionic bond? Which elements form ionic bonds? What is a covalent bond? Which two elements form a covalent (or polar covalent) bond? 12) What is electronegativity? What is the difference between a nonpolar and a polar covalent bond? 13) Under what conditions will ionic substances conduct electricity? Under what conditions will covalent substances conduct electricity? 14) Describe metallic bonding. Why are metals very good conductors of heat and electricity? 15) Draw Lewis structures for the following molecules: NH3, CO2, CO and BF3. Which of the four is exception to the octet rule and why?

Unit 10 Bonding~VSEPR, Dipoles/IMF’s 16) State the geometry (shape) and bond angles for the four molecules in problem #12. 17) Which of the 4 molecules in problem #12 are polar? Draw a net dipole arrow on the polar molecules. 18) What conditions are necessary for a hydrogen bond to occur between two molecules? 19) Why is water less dense in the solid state than in the liquid state? Include in your answer a reference to the intermolecular forces between the water molecules.

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Unit 11 Thermochemistry 20) How much energy is required to heat 50.0 g of lead (specific heat capacity = 0.128 J/g oC) from 25.0 oC to 500.0 oC? 21) A thick, clear liquid is poured onto a pile of purple crystals. Smoke is released, followed by a purple flame and intense heat. Is the reaction endothermic or exothermic? How does the energy of the reactants compare to the energy of the products? 22) Breaking bonds ______energy while forming bonds ______energy. 23) The temperatures of equal masses of granite and water are increased by 1 oC. The specific heat capacity of granite is 0.803 J/g oC. The specific heat capacity of water is 4.18 J/g oC. Which substance will have stored more heat energy (Q) and why? 24) Using the thermite reaction: 2 Al + Fe2O3 → Al2O3 + 2 Fe ∆H = -850 kJ a) Is this reaction endothermic or exothermic? Why? Is heat released or absorbed by the system? b) Calculate the heat released when 5.00 g of Al reacts with excess Fe2O3 pH mini-unit 1) Label each of the following solutions as being either acidic, neutral or basic: a) [H+] > [OH-] b) [H+] = [OH-] c) [H+] < [OH-] 2) For each of the following solution, calculate the pH, then label as acidic, neutral or basic: a) [H+] = 1.0 x 10-7 M b) [H+] = 1.0 x 10-8 M c) [H+] = 1.0 x 10-3 M

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Extra Credit Opportunity for Spring Final

For each of the Units, you may earn extra credit points by doing the following:

• WRITE OR TYPE 1 PAGE SUMMARY OF KEY IDEAS OF UNIT → 1 pt./ UNIT Include examples of solved problems as appropriate; No credit for handing in the notes/ materials already provided in workbook- must be a summary of your notes, not a copy. • SOLVE ALL OF THE PROBLEMS FOR A UNIT GIVEN BELOW → 2 pts/UNIT Honest effort, work must be shown for credit.

Total points possible = 5 pts (summaries) + 10 pts (problems) = 15 pts. Do not have to complete all units to earn points, however each subsection of work must be complete. For example, 1 pt earned for summarizing Unit 1, 2 pts earned for problems Unit 6 = 3 pts. ½ complete summary, ¾ problems attempted = NO CREDIT

Extra Credit Problems Posted On-line

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