UNIT 4a Thermochemistry Objectives • Student will understand energy and its forms, including kinetic, potential, chemical,

and thermal energies • Student will understand the law of conservation of energy and the processes of heat

transfer • Student will use thermochemical equations to calculate energy changes that occur in

chemical reactions and classify reactions as exothermic or endothermic. • Student will perform calculations involving heat, mass, temperature change, and

specific heat • Student will use calorimetry to calculate the heat of a chemical process.

UNIT 4a VOCABULARY

• Heat (q) • System • Surroundings • Thermal equilibrium • Calorimetry/Calorimeter • Calorie (cal) • (Dietary) Calorie (Cal) • Kilocalorie • Joule (J) • Specific heat (capacity)

• Thermochemical equation • Enthalpy (H)/Heat of reaction (ΔH) • Entropy (S)/ Entropy of reaction

(ΔS) • Gibbs Free Energy (G)/Free energy

change (ΔG) • Latent heat • Heating curve • Molar enthalpy/heat of

o fusion (ΔHfus)

Lesson Topic pg

Unit 4a Unit 4a - Thermochemistry 1-2

4.1a Intro Thermo Notes 2-3

4.2a Measuring Heat Calorimeter Problems 3-6

4.3a Hess’s Law and Heat of Reactions 6-8

4.4a Entropy/Gibbs Free Energy 9-11

4.5a Hand Warmer & Cold Pack Thermo Lab 11-15

4.6a Catch-up Day

4.7a Exam

Unit 4b Unit 4b – Nuclear Chemistry 16

4.8b Introduction to Nuclear Chemistry 16-20

4.9b Nuclear Chemistry Practice Questions 20

4.10b Nuclear Chemistry Projects 21-24

4.11b Nuclear Chemistry Project Workday 21-24

4.12b Catch-up Day

4.13b Exam

Homework Problems 24-30

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o solidification (ΔHsolid) o vapozation (ΔHvap)

o condensation (ΔHcond)

• Standard heat of formation (ΔHf

°) • Hess’s law • Exothermic • Endothermic

• Energy • Work • Change of state • Solid • Liquid

Lesson 4.1 a Thermochemistry: heat changes that occur during chemical reactions

Define Energy:

In the context of this chemistry course, this means that work is done when an object moves some distance in response to a force being applied. Work = force x distance.

Types of energy • Mechanical Energy

o Potential or stored energy ▪ Wound up spring ▪ Gasoline before burning

• Kinetic energy or energy of motion o Moving objects

• Nonmechanical energy o Chemical energy o Electrical energy o Electromagnetic (radiant) energy o Sound energy o Heat energy o Magnetic energy

Chemical Potential Energy: Energy is often converted from one form to another. In a car the chemical energy of the fuel is converted to heat energy, which is then converted to mechanical energy as gases expand to force the pistons to move.

Law of Conservation of Mass/Energy As with the conservation of mass, energy cannot be either created or destroyed in a reaction. Einstein’s research led to the discovery that mass can be converted into energy (E=mc2)

Heat (q): energy that transfers from one object to another based on temperature differences. Heat always moves from hot to cold or high average kinetic energy to low average kinetic energy

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System:

Surroundings:

Law of Conservation of Energy: energy is never created nor destroyed it can all be accounted for in terms of work, stored energy, or heat Endothermic: system absorbs energy from surroundings (+ q value). Energy (heat) is absorbed as the products of the reaction have more potential energy than reactants. (Heat + NH4Cl + H2O ! NH4OH + HCl)

Examples • Photosynthesis • Melting and evaporation

Exothermic: system releases heat to surroundings ( - q value). Energy (heat) is released as the products of the reaction have less potential energy than reactants. (CH4 + O2 ➔ CO2 + 2H2O + Heat).

Examples • Burning • Rusting of iron • Condensation • Freezing

Units for measuring energy Joule (J) Calorie (cal)

calorie: quantity of heat needed to raise the temperature of 1 g (1 ml or 1 cc or 1 cm3) of pure H2O 1 C°

1 Calorie (dietary calorie) = 1 kilocalorie = 1000 calories

SI unit for heat and energy is Joule (J) named after James Prescot Joule

1 J = 0.2390 cal or 4.184 J = I cal

Heat capacity (C): amount of energy needed to raise or lower a substance 1°C. Heat capacity depends on makeup of substance as well as mass of substance

I mL of water takes less heat than 100 ml of water to raise I degree C

Specific Heat Capacity (c): the amount of heat energy needed to raise or lower the temperature of 1 g of a substance by 1 oC.

4.2a Measuring Heat Calorimeter Problems Heat energy- Each substance has its own specific heat capacity which is defined as the amount of heat required to raise one mass unit 1 oC (J/g-°C or cal/g-°C) c = specific heat

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c = heat (q) (ioules or calories) rearranged q = m c ∆T mass(g) x ∆T (oC)

Kinetic theory of Heat and Temperature All matter is made of uncountable numbers of particles. All particles are in constant motion. Even solid objects are in motion, even though we cannot detect motion. With two bodies in contact, energy will transfer from the body with the higher energy to the body with the smaller amount of energy until the temperature is the same. Melting of ice is an example of this transfer. Heat enters melting ice, but as long as both solid and liquid are present the temperature does not change. The energy is being used to rearrange the molecules. The potential energy of the liquid water is higher than the ice, therefore the melting of ice is endothermic. Evaporation of sweat is another example.

Specific Heat and Calorimetry q=mcΔT 1. How much heat is required to raise the temp of 654 g of water from 34.5°C to 89.7°C?

Ans ______ 2. How much heat is required to raise the temp of 654 g of silver from 34.5°C to 89.7°C?

Ans ______

More Thermodynamics:

Remember: If q > 0 we say that the object gained heat - heat flowed into our sample (endothermic), the system absorbed heat. Likewise if q < 0 we say that our object lost heat, heat flowed out of our sample (exothermic). The system lost heat.

The temperature change, ∆T, is also related to heat. It can be positive or negative. It is obvious that if ∆T> 0 then the temperature increased and if ∆T < 0 then the temperature decreased.

Clearly whenever ∆T> 0 then q > 0, and vice versa.

Changes of State: If you have a glass of ice water (the glass contains both ice and liquid water) and you add some heat to the ice water the temperature does not change. The heat melts some the ice rather than changing the temperature of the system. Fill in the phase change diagram for H2O (below):

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substance c (J/g°C) water 4.184 ethanol 2.452 graphite 0.720 diamond 0.502 iron 0.444 copper 0.385 silver 0.237 gold 0.129

The conversion of ice to water (at 0 °C) is called a “phase change.”

Likewise, if you added some heat to a beaker of water at 100°C (the normal boiling point of water) the temperature of the water would not increase.

You would just vaporize some of the water. The conversion of liquid water to water vapor is another phase change. There is no increase or decrease in temperate when heat is added or withdrawn at a phase change.

Phase changes require energy (in the form of heat). You have to add heat to melt ice or to vaporize water.

Once again, the amount of heat required to make a phase change depends on the material and the amount of material.

There are tables of heats of fusion (sometimes called latent heat of fusion) and heats of vaporization (or latent heat of vaporization).

The heats of phase changes are usually listed in the tables in J/mol, although you can calculate the value in J/g by dividing by the molar mass of the element, compound, or molecule.

Specific Heats of H2O water 1.00 cal/g°C ice 0.50 cal/g°C steam 0.48 cal/g°C mp = 0.0°C (273 K) bp = 100.0°C (373 K)

Heat of fusion: Hf = 80.0 cal/g (6.008 KJ/mol or 6008 J/mol), (333.4 J/g).

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Heat of vaporization: Hv = 540. cal/g (373 K) (40.66 KJ/mol or 40660 J/mol), (2256 J/g)

Sample calculations:

1) How much heat (in cal) does it take to vaporize 1.000 L of H2O at 100°C? (1 mL of H20 = 1 g)

2) How much heat (in J) is released when 10.87 mL of liquid H20 freezes at 0°C?

4.3a Hess Law and Heat of Reactions

Enthalpy (ΔH)

Enthalpy is defined as the quantity of heat transferred in a reaction. The equation ΔHreaction = Hproducts - Hreactants illustrates how heat is calculated. The coefficients in a thermochemical reaction indicate the number of moles involved. It is possible to use fractions as coefficients in this kind of equation. In an exothermic reaction the amount of heat released is included in the equation:

2H2 + O2 → 2H2O + 483.6kJ. In an endothermic reaction heat is one of the reactants: 2H2O + 483.6kJ → 2H2 + O2 halving or doubling the amounts will change the amount of heat transferred correspondingly.

Molar Enthalpy of Formation (ΔHf) The enthalphy change that occurs when one mole of a compound is formed at standard state (25°C and 1atm). Compounds that release a great deal of heat upon formation are very stable. Defined in terms of one mole of product.

Molar Enthalpy of Combustion (ΔHc) The enthalpy change that occurs during the complete combustion of 1 mole of a substance. Defined in terms of one mole of reactant.

Standard values of ΔHf and ΔHc are available in many reference books. To calculate overall ΔH, use Hess’s Law, which states that the overall enthalpy change in a reaction is equal to the sum of the enthalpy changes in the individual steps in the reaction, independently of the route taken. This approach can help calculate the heat transferred when standard amounts are unavailable.

1. If a reaction is reversed, the sign of ΔH must be reversed. 2. Multiply the coefficients of the known equations so that when added together they give the desired thermochemical equation. Multiply the ΔH by the same factor.

ΔH = sum of [(ΔHf of products) x (mol of products)] –

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sum of [ΔHf of reactants) x (mol of reactants)]

Heat of Reaction 1. What is meant by enthalpy?

List three different types of enthalpies

2. What is the general equation to calculate enthalpy change?

3. ΔH will always be positive for an (endothermic/exothermic) reaction, and ΔH will always be negative for an (endothermic/exothermic) reaction.

4. What is meant by heat of formation?

5. How many kJ are released when 1 mole of Pb(NO3)2 is formed from its elements? (see reference sheet): ans____________

6. Given the following reaction: Si(s) + O2 (g) ---> SiO2(s) ΔHf= -910.7 kJ/mol

Calculate ΔHf for 2 moles of SiO2: ______ Calculate ΔHf for 0.5 moles of SiO2: _____

Hess’s Law Introduction Recap Enthalpy change (represented by ΔH) is the difference in enthalpies of reactants and products during a change.

Enthalpy means “thermal energy”.

Hess’s Law Hess’s law states that ΔH (total enthalpy change) is the sum of all intermediate steps taken to arrive at the target. It is a mathematical relationship which makes a connection between a series of reactions which change a group of reactants into a group of products.

This can be written in a formula as such: ΔHtarget = ΣΔHunknown

Hess’s law can also be expressed by saying that if two or more equations for which the enthalpy is given or otherwise known can be added together to create a target equation, the enthalpy changes can also be summed to find the enthalpy change of the target equation.

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The sum of

HESS’S LAW PRACTICE PROBLEMS

1) Calculate ΔHo for the formation of 1 mol of strontium carbonate (the material that gives the red color in fireworks) from its elements.

2) The combination of coal and steam produces a mixture called coal gas, which can be used as a fuel or as a starting material for other reactions. If we assume coal can be represented by graphite, the equation for the production of coal gas is:

3) Find the ΔH for the reaction below, given the following steps and subsequent ΔH values: N2H4(l)  +  H2(g)  →  2NH3(g)

N2H4(l)  +  CH4O(l)  →  CH2O(g)  +  N2(g)  +  3H2 (g)         ΔH = -37 kJ N2(g)  +  3H2(g)  →  2NH 3(g)                                                ΔH = -46 kJ CH4O(l)  →  CH2O(g) +  H 2(g)                                              ΔH = -65 kJ

4.4a Entropy and Gibbs Free Energy

The second law of thermodynamics states that there is an inherent direction in which a reaction, which is not at equilibrium, moves. If this direction is to the right (toward the products) the reaction is said to be spontaneous. Reactions which occur under normal conditions of temperature and pressure without assistance are called spontaneous reactions. Reactions which are spontaneous in one direction are nonspontaneous in the reverse direction.

Entropy is the thermodynamic function which measures the degree of randomness of a chemical system. Entropy is given the symbol, S. If the change in entropy during a chemical reaction, ∆S, is positive, this indicates an increase in randomness or disorder. If ∆S is negative, this indicates a decrease in randomness.

kJ 206.1 H (g) CO (g) H 3 (g) OH (g) CH (3)kJ 41.2 - H (g) H (g) CO (g) OH (g) CO (2)kJ 131.3 H (g) H (g) CO (g) OH C(s) (1)

:reaction of enthalpies standard following thefromreaction for this changeenthalpy standard theDetermine

(g)CO (g) CH (g) OH 2 (s) C 2

224

222

22

242

=°Δ+→+

=°Δ+→+

=°Δ+→+

+→+

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3 2 32

1 2 2

2 3

Sr (s) C(graphite) O (g) SrCO (s) The information available is(1) Sr (s) O (g) SrO (s) H -592 kJ(2) SrO (s) CO (g) SrCO

+ + →

+ → Δ ° =

+ →

2 2

(s) H -234 kJ(3) C(graphite) O (g) CO (g) H -394 kJ

Δ ° =

+ → Δ ° =

The second law of thermodynamics is stated as: the entropy of the universe will increase for any spontaneous process.

∆Suniverse = ∆Ssystem + ∆Ssurroundings

For a spontaneous process, ∆S of the universe always increases.

Now, a man named Gibbs became very famous by renaming the left part of the above equation after himself, and giving it a new symbol.

WHEN ΔG < 0 , the process DISORDERS the universe is SPONTANEOUS WHEN ΔG > 0, the process ORDERS the universe is NON spontaneous WHEN ΔG = 0, the process is at EQUILIBRIUM

∆G is known as Gibb's free energy, and it is the maximum amount of useful work available from a spontaneous reaction, calculated by taking the amount of useful energy from a reaction (heat) and subtracting out the amount of un-useful energy from a reaction (entropy or disorder). It is also the minimum amount of work required to cause a nonspontaneous reaction to occur. When ∆G˚ for a reaction is negative, the reaction is spontaneous; if ∆G˚ is positive, the reaction is nonspontaneous. This is because of the negative sign, so the rules are opposite. We rewrite the equation above as:

∆G˚ = ∆H˚ - T∆S˚

So ∆G˚ for a chemical reaction can be determined knowing ∆H˚ and ∆S˚ for the chemical reaction. Notice that a negative enthalpy (exothermic) helps a reaction become spontaneous, and a positive entropy (increasing disorder) helps a reaction become spontaneous.

∆G˚ and ∆S˚ values are also tabulated in the back of the book, and Hess's Las applies to those as well.

∆G = Σ ∆G- Σ ∆G ∆S = Σ S- Σ S

The following table show the temperature dependence of a reaction on temperature. To determine how a reaction will respond to a temperature change, you must take both enthalpy and entropy into account. By examining the sign of enthalpy and entropy and using the equation:

∆G˚ = ∆H˚ - T∆S˚

it can be determined what range of temperatures a reaction will be spontaneous at.

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The magnitude of both ∆H˚ and ∆S˚ change with temperature. However, it is assumed the amount of the change is small over small temperature changes.

Practice Problems:

1. Predict whether the entropy of the system increases, remains constant or decreases when the following processes occur. Explain your reasoning. a) Ice melts at 0 °C.

b) A precipitate forms in aqueous solution.

c) A solid dissolves in water.

d) A gas condenses to a liquid.

2. Calculate ∆S°reaction, ∆H°reaction and ∆G°reaction (both ways) for the following chemical reaction at STP.

Sign of ∆H˚rxn (25 ˚C)

Sign of

∆S˚rxn (25 ˚C)Sign of

∆G˚rxn (25 ˚C)

Sign of ∆G˚rxn at high

temperature

Sign of ∆G˚rxn at low

temperature

+ + – or + ∆Grxn is - at high T

∆Grxn is + at low T

+ – + ∆Grxn is always + ∆Grxn is always +

– – + or – ∆Grxn is + at high T

∆Grxn is - at low T

– + – ∆Grxn is always - ∆Grxn is always -

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Al(s) + Cl2(l) → AlCl3(s)

4.5a Designing a Hand Warmer & Cold Pack Lab

Background: Hand warmers are familiar cold weather gear used to quickly provide warmth to frigid fingers. Many commercial hand warmers consist of a plastic package containing a solid and an inner pouch filler with water. When the pack is activated, the solid dissolves in water and produces a large temperature change. The energy or enthalpy change associated the process of a solute dissolving in a solvent is called the heat of solution (ΔH soln). At constant pressure, this enthalpy change, ΔHoln, is equal in magnitude to the heat loss or gain, q, to the surroundings. In the case of an ionic solid dissolving in water, the overall energy change is the net result of three processes- the energy required to break the attractive forces between ions in the crystal lattice (ΔH1 = +C kJ/mol), the energy required to disrupt intermolecular forces between water molecules ΔH2 = +D kJ/mol), and the energy released when the dissociated (free) ions form ion-dipole attractive forces with the water molecules )( ΔH3 = -F kJ/mol). The overall process can be represented by the following equation.’

MaXb(s) ! aM+b (aq) + bX–a (aq) ΔHsoln, = ΔH1 + ΔH2 + ΔH3 = (+C + D + F)kJ/mol

ΔSoln, = S1 + + S2 + S3 = (+C + D + F)kJ/mol If the amount of energy released in the formation of hydrated ions (ΔH3) is

greater than the amount of energy required to separate the solute and solvent particles (ΔH1 + ΔH2), then the sum (ΔHoln,) of the energy changes will be negative and the solution process exothermic (releases heat). If the amount of energy released in the formation of hydrated ions is less than the amount of energy required to separate the solute and solvent particles, then the sum of the energy changes will be positive and the solution process endothermic (absorbs heat). What would the sign of ΔSsoln ?

Heat of solution and other enthalpy changes are generally measured in an insulated vessel called a calorimeter that reduces or prevents heat loss to the atmosphere outside the reaction vessel. The process of a solute dissolving in water may either release heat into the resulting aqueous solution or absorb heat from the solution, but the amount of heat exchanged between the calorimeter and the outside surroundings should be minimal. When using a calorimeter, the reagents being studied are mixed directly in the calorimeter and the temperature is recorded both before and after the reaction has occurred. The amount of heat transfer (q) may be calculated using the heat energy equation: q – m x s x ΔT (Equation 1). Where m is the total

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mass of the solution (solute + solvent). The specific heat of the solution is generally assumed to be the same as that of water, 4.18 J/g oC.

When measuring the heat transfer for an exothermic heat of solution using a calorimeter, most of the heat released is absorbed by the aqueous solution (qaq). A small amount of heat will be absorbed by the calorimeter itself (qcal). The overall heat transfer (qsoln) for the reaction (the system) then becomes: qsoln = qaq + q cal (Equation 2). In our experiments we will assume that the qcal is negligible.

Experimental Overview The purpose of this lab is to design an effective hand warmer that is inexpensive, nontoxic, and safe for the environment.

Material Safety Data Sheets (MSDS)

Safety Precautions Ammonium nitrate is a strong oxidizer and may explode if heated under confinement. It is also slightly toxic by ingestion and a body tissue irritant. Calcium chloride is slightly toxic. Lithium chloride is moderately toxic by ingestion. Magnesium sulfate is a body tissue irritant. Sodium acetate is a body tissue and respiratory tract irritant. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves, and a chemical apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Part I: Calorimetry PRACTICE Procedure

Examine the heat energy change for the following solution. MgSO4(s) + H2O(l) ! Mg2+ (aq) + SO42- (aq)

1. Measure 45.0 mL of deionized water in a graduated cylinder and transfer to calorimeter. 2. Measure and record the initial temperature of the water. 3. Measure 5.00 g of anhydrous magnesium sulfate (on piece of paper or weigh boat) 4. Put a magnetic stir bar or use stirring rod into calorimeter and slowly stir the water. Again

if available not necessary just make sure the solution mixes by swilling the cup. 5. Quickly add the 5.00 g of magnesium sulfate to the calorimeter and insert the

thermometer. 6. Monitor the temperature and record the highest (or lowest) temperature reading. 7. Calculate the molar heat of solution for magnesium sulfate. Include the correction due to

the heat capacity of the calorimeter. Data Table

Solid NH4NO3 CaCl2 KCl NH4Cl Na2CO3 NaNO3

Cost ($)/kilogram

18.50 10.80 30.00 1.67 5.95 20.00

Mass of Water Trial I Trial 2

Mass of MgSO4(s)

Total Mass (Water + MgSO4(s))

Initial Temperature

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Part II: Design a Hand Warmer or Cold Pack Challenge You are challenged to use chemistry to design an effective, safe, environmentally benign, and inexpensive hand warmer or Cold Pack. The ideal hand warmer increases in temperature by 20oC (but no more) as quickly as possible, has a volume of about 50 mL and uses 10 grams of an ionic solid, costs as little as possible to make, and uses chemicals that are as safe and environmentally friendly as possible. The same argument can be made for a chemical Cold Pack.

Each group will use 3 solids and then collaborate with another group. 1. Review the calorimetry procedure and answer the following questions in your course

packet or lab handout. a. What data is needed to calculate the enthalpy change for a reaction?

b. Identify the variables that will influence the experimental data

c. What variable should be controlled (kept constant) during the procedure?

d. The independent variable in an experiment is the variable that is changed by the experimenter, while the dependent variable responds to or depends on the changes in the independent variable. Name the independent and dependent variables in a calorimetry experiment to determine the molar heat of solution.

e. Discuss the factors that will affect the precision of the experimental results.

Data Table of 6 Ionic Compounds (pick two and run three times)

Final Temperature

qMgSO4

Calculated ΔHsoln

Salt T1: Salt(1)

T2: Salt(1)

T3: Salt(1)

T1: Salt(2)

T2: Salt(2)

T3: Salt(2)

Mass of Water

Mass of Salt

Total Mass (water+salt)

Initial Temperature

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Questions: 1. Write a balanced equation for the reaction you studied (including the heat).

2. Dissolving ionic compounds involves the separation of the solid ionic compound into cations and anions in water. This process can be represented by an equation showing the solid as a reactant and the aqueous ions as products. The heat of reaction, ΔHsoln, is written after the products, typically in units of kJ/mol. Example: sodium hydroxide dissolves exothermically, releasing 44.2 kJ/mol dissolved. This process is represented as: NaOH(s) ! Na+(aq) + OH-(aq), ΔHsoln = -44.2 kJ/mol. Write an equation to represent the dissolving process for each salt you studied. Include your calculated heat of reaction as in the example.

3. Was the reaction spontaneous? How do you know this?

4. From the temperature change of your trials, what must be the sign for ΔH?

5. From question 3, what must be true about the sign for ΔS? Why?

6. What are the units for entropy, ΔS?

7. Find the published value of ΔHsoln for each solid and determine the % error in the class average value.

Final Temperature

Calculated qsalt

Calculate ΔHsoln

Calculate ΔSsoln

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8. What possible sources of error could affect the accuracy of your calculated value of the amount of sold in your hand warmer or cold pack? List at least two and what effect they would have on the temperature change.

Lesson 4b Nuclear Chemistry Objectives • Distinguish between alpha, beta, and gamma radiation • Explain the cause of spontaneous decay • Determine the decay mode of a radioactive isotope • Explain the law of conservation of mass • Write nuclear equations showing alpha and beta decay • Complete nuclear equations by predicting missing particles • Interpret radioactive decay series graphs • Define half-life • Calculate half-life • Distinguish between natural and artificial nuclear reaction • Distinguish between fusion and fission nuclear reaction

Unit Vocabulary • Radioactive isotope • Spontaneous decay • Radioactive particle • Alpha Particle • Beta Particle • Gamma Ray • Transmutation (natural and artificial) • Half-life • Fusion • Fission

4.8b Introduction to Nuclear Chemistry

!

Outline• Background/Introduction• Spontaneous Nuclear Reactions

- Types of Radioactive Decay- Radioactive Half-Life

• Stimulated Nuclear Reactions- Nuclear Fission- Nuclear Fusion

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Nuclear Chemistry - Background• Traditional Chemistry– Reactions occur due to interactions between valence

electrons (surrounding nucleus)

• Late 1800s – Early 1900s ! New Developments– Discovery that Uranium emits radiation (Henry Becquerel)– Amount of radiation emitted is proportional to amount of

element present (Marie Curie)– “Radioactive” substances = radiation-emitting (Curie)– Radioactivity = a inherent property of certain ATOMS, as

opposed to a chemical property of compounds (Curie)

• Birth of nuclear chemistry

Chemical vs. Nuclear ReactionsChemical Reactions Nuclear Reactions

• Loss, gain, sharing of valence electrons

• Formation of compounds(Recall the different reaction types)

• Affected by temperature, pressure,presence of other atoms

•Changes in nucleus

• Transformation of one elementinto a different isotope or adifferent element altogether

•Unaffected by temperature,pressure, presence of other atoms

Which Elements/Isotopes Are Radioactive?• Based on the nuclear force and nuclear stability

–More protons in nucleus ! more neutrons needed to bind nucleus together

– Critical Factor = Neutral-to-Proton Ratio• Neutron-to-Proton ratios of stable nuclei increase with

increasing atomic number• Unstable neutron-to-proton ratio = RADIOACTIVE

– Nuclei with atomic numbers ≥ 84 = ALWAYS Radioactive• Very large nuclei!• Neutron-to-proton ratio always unstable

Nuclear Stability• How can a stable nucleus exist?– Net positive charge in nucleus surrounded by negative

charges (electrons)– Electrostatic force

• Opposite charges attract• Like charges repel

– Why doesn’t the nucleus break apart?

• Nuclear Force (1934)– Force between protons and neutrons than binds nucleus together

within atom– Very strong (must be!)

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Radiation and Radioactive Decay• Unstable (radioactive) nuclei emit radiation (energy) in

order to become more stable

• Radioactive Decay – occurs when a nucleus spontaneously decomposes in this way

• 3 Common Types of Nuclear Reactions– Alpha Radiation– Beta Radiation– Gamma Radiation

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Half-Life Problems• Can be completed without the below formulas!– Conceptual knowledge of half-life is sufficient

Half-Life Formulas

Fraction remaining = (n = # of half-lives elapsed)

Amount remaining = Original Amount * Fraction Remaining

n21

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Lesson 4.9b Nuclear Chemistry Practice Questions

19

1. Complete the following table:

2. What is the atomic weight of silicon, if 92.2% of its atoms have mass 28.0 g/mol, 4.7% have mass 29.0 g/mol, and 3.1% have mass 30.0 g/mol?

Atomic weight__________________

3. Complete the following table:

4.10 &11b Nuclear Energy Timeline Project Option 1 For this assignment you will create a timeline of nuclear energy from its

discovery to the present. You will be given a large sheet of paper to prepare your time line. Be sure that you decide upon the correct spacing of your years so that you use the entire piece of paper. You will be given two class days to work on the timeline in class.

You will be required to find facts and dates about the different topics. Any important dates and what occurred must be added as a minimum requirement, however you may add any more information that you feel is important. As part of the time line you may wish to include pictures of people, maps, or buildings.

The internet is a wonderful tool however it is very easy to cut and paste information. Any group caught plagiarizing will automatically receive a zero on the assignment and will fail the semester.

Your time line must include the following points: Marie Curie

- When was she born

Isotopic Notation # of Protons

# of Neutrons

Mass Number

Element Name

a. 12 H

b. 1737Cl

c. Silver-107

Charge Penetrating Ability

Alpha (24 He)

Gamma (γ)

Beta ( ) e01−

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- What did she win her Nobel prizes for - What was her great nuclear energy discovery - When did she die and what of?

Albert Einstein - When was he born? - When did he die?

- When did he discover the Theory of Relativity - What was his contribution to the discovery of radiation

Discovery of radiation - What was used to detect the radiation - Three main types of radiation and when and who discovered them

Enrico Fermi - What did Fermi win his Nobel Prize for - What was his great discovery

Discovery of X-ray - Who discovered it First self-sustaining chain reaction

- Who were the members of the team - Where did the project take place - What element was used

Manhattan Project - who were the members of the project - Where was the project conducted - What was significant about the project

Production of the first nuclear weapon - Where was the first bomb tested - What technology is used in this nuclear weapon - how many countries now have nuclear weapons capabilities (add these

countries based on date of development to your timeline) Hiroshima

- What was the name of the bomb dropped on the city - How many people were killed - What was the area of destruction - Why was the city chosen - What is the condition of the city today

Nagasaki - What was the name of the bomb dropped on the city - How many people were killed - What was the area of destruction - What is the condition of the city today

Development of the first Nuclear Power Plant - Where was the first power plant built

o What was first powered by the plant - How many nuclear power plants are there in the US - How many countries use nuclear power as a form of energy

Last nuclear power plant built in the US - What is its name - Where is it located - Is it still in use today

Three Mile Island - What is Three Mile Island - Where is Three Mile Island - What happened at Three Mile Island

o What caused the problem 21

- Were there any deaths - Is Three Mile Island still in use today - Is the area contaminated

Chernobyl - Where is Chernobyl - What is Chernobyl - What happened at Chernobyl

o What caused the problem - Were there any deaths - Is Chernobyl still in use today - What is the aftermath of the area (are people developing illnesses, is the

area contaminated, is the land useable….) Fission

- Who discovered it - Where was it discovered - What are benefits of fission - What are the problems with fission - What are the uses of fission

Fusion - Who discovered it - Where was it discovered - What are benefits of fusion - What are the problems with fusion - What are the uses of fusion

Atomic Energy Commission (AEC) - Why was it created and what is its purpose

First Nuclear Submarine - What is its name

International Atomic Energy Agency (IAEA) - Why was it created and what is its purpose

Treaty for Non-Proliferation of Nuclear Weapons - Who was the first to sign - what is its purpose - Who are the five recognized nuclear powers

Fukushima Nuclear Disaster - When did it occur - What caused it - How was it changed attitudes about nuclear power in Japan

Today (2015) - How many nuclear power plants are in use in the US - What percent of energy do the produce

1st and last Texas nuclear power plant developed - What are their names - Where are they located - When did construction first begin - When did construction end

4.10 &11b Nuclear Energy Flyer Project Option 2 Project Description For this project you are going to assume that an action group is trying to lobby the city council of Austin to ban the use of ALL nuclear substances in the Austin area. You will create a flyer to present the facts of nuclear chemistry, both the good and the bad, using scientific evidence. To do this, you need to describe the characteristics of

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alpha, beta, and gamma radiation. In this description, be sure to describe the role of the decay process using balanced nuclear reactions and compare and contrast the use of fission and fusion in nuclear chemistry. Evaluate the impacts this ban will have on the citizens of Austin. Essential Questions 1. How can radiation be harmful and helpful? 2. Do the risks of nuclear energy outweigh the benefits Vocabularynon-ionizing radiation, ionizing radiation, radioactive decay, alpha particle, beta particle gamma ray, nucleon, radioisotopes radioactive decay, transmutation, fission , fusion chain reaction, half-life

Objectives 1. Create a flyer to evaluate the good and bad of nuclear chemistry using scientific

evidence. 2. Describe the characteristics of alpha, beta, and gamma radiation. 3. Describe the radioactive decay process using balanced chemical equations. Compare fusion and fission reactions.

Homework Practice Problems: Lesson 4.1a Introduction to Thermochemistry

1. If 7350 J were added to 152 g of ethanol, its temp would go up by how much?

Ans _______ 2. 16.25 g of water at 54.0°C releases 402.7 J. What will be

its final temp?

Ans _______ 3. 697 J are added to a 36.8 g of kerosene and the temp

increases from 22.5°C to 34.7°C. Determine kerosene's specific heat.

Ans ______

4. 25 copper pennies (each weighing 3.12 g) are placed in 36.0 g of ethanol at room temp (22.1°C). How much heat will it take to raise the temperature up to 65.8°C?

Ans _______

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substance c (J/g°C) water 4.184 ethanol 2.452 graphite 0.720 diamond 0.502 iron 0.444 copper 0.385 silver 0.237 gold 0.129

5. What mass of 54.0°C water must be added to 468 g of 21.0°C water to make the final temp of both come out to be 29.0°C?

Ans _______ 6. What mass of 54.0°C gold must be added to 468 g of 21.0°C water to

make the final temp of both come out to be 29.0°C?

Ans _______

7. A 325 g brass rod at 100.0°C is placed in a cup containing 162 g of 24.3°C water. The final temp comes out to be 37.4°C. Determine brass's specific heat.

Ans _______ 8. 100.0 g of water at 20.0°C are mixed with 200.0 g of copper at 40.0°C.

What will the final temp come out to be?

Ans _______

Lesson 4.2a Thermochemistry Phase Changes

1. How much heat (in joules) is required to change 45.40 g of ice at 0.0°C into water and then raise the water temp to 23.0°C

2. How much heat (in cal) is given off when 18.5 g of water at 35.0°C freezes to ice and then is lowered to -5.00°C?

Lesson 4.3a Hess Law

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1. What is ΔHf for this reaction? 2 P(s) + 5/2 O2 (g) ---> P2O5 (s)

2. What is ΔHf for this reaction? 4 P(s) + 5 O2 (g) ---> 2 P2O5 (s)

3. How many kJ are released when 1 mole of beer alcohol (C2H5OH) is burned?

4. Calculate the heat of formation for the combustion of one mol of ethanol.

5. What is the value for ΔHc when 249.2 g of sugar (sucrose) is burned?

6. Find the ΔH for the reaction below, given the following steps and subsequent ΔH values: H2SO4(l)  →  SO3(g)  +  H2O(g)

a. H2S(g)  +  2O2(g)  →  H2SO4(l)                                  ΔH = -235.5 kJ H2S(g)  +  2O2(g)  →  SO 3(g)  +  H2O(l)                    ΔH = -207 kJ H2O(l)  →  H2O(g)                                                      ΔH = 44 kJ

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7. Find the ΔH for the reaction below, given the following steps and subsequent ΔH values: ½ H2(g)  +  ½ Cl2(g)  →  HCl(g)

a. COCl2(g)  +  H2O(l)  →  CH2Cl2(l)  +  O2(g)                               ΔH = 47.5 kJ 2HCl(g)  +  ½ O2(g)  →  H 2 O(l)  +  Cl2(g)                                 ΔH = 105 kJ CH2Cl2(l) +  H2(g)  +  3/2 O 2(g)  →  COCl2(g)  +  2H 2O(l)         ΔH = -402.5 kJ

Lesson 4.4a Entropy and Gibbs Free Energy

1. Consider the following reaction:

N2(g) + 3H2(g) --> 2NH3(g) The reaction is carried out at 100oC

a. Calculate the change in entropy of the system.

b. Calculate the change in enthalpy, ΔHreaction.

c. Calculate the change in Gibbs free energy, ΔGreaction at 100oC.

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d. Is the reaction spontaneous?

2. The enthalpy of combustion, ∆H°comb, for oxalic acid, C2H2O4(s), is -246.05 and

i. Substance ∆H°f S° ii. C(s) 0 5.69 iii.CO2(g) -393.5 213.6 iv. H2(g) 0 130.6 v. H2O(l) -285.8 69.96 vi.O2(g) 0 205

vii.C2H2O4(s) → 120.1

b. Write the balanced chemical equation that describes the combustion of one mole of oxalic acid.

c. Write the balanced chemical equation that describes the standard formation of oxalic acid.

d. Using the information given above and the equations in a) and b), calculate ∆H°f for oxalic acid.

e. Calculate ∆S°f for oxalic acid and ∆S°rxn for the combustion of one mole of oxalic acid.

f. Calculate ∆G°f for oxalic acid and ∆G°rxn for the combustion of one mole of oxalic acid.

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g. Is the formation of oxalic acid from its elements spontaneous? Is the combustion of oxalic acid at 25 °C spontaneous?

Lesson 4.8b Nuclear Chemistry Practice Problems

1. Complete the following: a. Alpha emission: 92235U →_________________________________________

b. Beta emission: 83210Bi→__________________________________________

c. Alpha emission 88213Ra→ __________________________________

2. Define half-life

3. The half-life of chromium-51 is 28 days. If you have 2.00 grams of chromium, how many grams would remain after 56 days?

4. The half life of Carbon-14 is 5730 yrs. What fraction of the original sample will still be Carbon-14 after 17190 yrs have passed?

5. In the year 1992, a doctor’s office purchased 2 grams of Cesium-137 that registered 400 radiation counts per second (cps). Given the half-life of Cesium is close to 30 yrs, how much radiation would be expected to remain by the year 2015?

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6. What type of particle is A____________, B_____________, C________________

7. Write a nuclear equation for the alpha decay of each of the below radioactive isotopes:

a. Radium-226 b. Thorium-230

8. Write a nuclear equation for the beta decay of each of the below radioactive isotopes:

a. Radon-231 b. Krypton-81

9. Radium-226 undergoes three consecutive episodes of alpha decay followed by two episode of beta decay, followed by an episode of alpha decay. What isotope is present after this nuclear activity?

10.Sodium-25 was to be used in an experiment, but it took 3.0 minutes to get the sodium from the reactor to the laboratory. If 5.0 mg of sodium-25 was removed from the reactor, how many mg of sodium-25 were placed in the reaction vessel 3.0 minutes later if the half-life of sodium-25 is 60 seconds?

11.A 2.5 gram sample of an isotope of strontium-90 was formed in a 1960 explosion of an atomic bomb at Johnson Island in the Pacific Test Site. The half-life of strontium-90 is 28 years. In what year will only 0.625 grams of this strontium-90 remain?

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