Electron Configuration Elements (atoms) and Ions · Web viewArgon has three naturally occurring isotopes: argon-36, argon-38, and argon-40. Based on argon’s reported atomic mass,

IB Chem I – Stations Review – Topic 2.2, 2.1

Name:_________________________________Partner Name:_____________________ Date:__________

Work through the stations below with your partner in any order you would like. Check your answers with the Station Keys. Then, answer the Station Mini Quiz at the end. Plan to spend ~10 minutes at each station. Work efficiently! Record your answers and create your OWN question to leave behind for other groups! But don’t leave them the answer. Extra points will be awarded to the best question and you might see it on the test…

The review packet at the end of the stations is to be completed outside of class for your own independent practice. Keys will be posted on the website and it will not be collected. Please complete the additional review packet by September 14th and bring any questions you may have to go over during class before the review game.

Keys can be found at http://shakeribchem.weebly.com

Go to Unit 1b – Topic 2 and scroll to the last links on the page.

Unit 1b Exam Topic 2.1 / 2.2: September 15th, 2017 ONLY – no extra time!

MC Section 20%

Short Answer & Problem Solving80%

ContentsPage #

Station 1 – Electron Configuration & The Rules2

Station 2 – Anomalies & Excited States4

Station 3 – Extension: Trouble with the Aufbau principle6

Station 4 – Electron Behavior, Energy, and Light7

Station 5 – Emission Spectra and Analysis9

Station 6 – Calculating Photon Energy11

Station 7 – The Bohr Model and Spectra12

Station 8 – Extension: Ionization Energy Predictions13

Station 9 – Average Atomic Mass and Isotopes14

Station 10 – Calculating Percentage Abundances16

Review Packet – Independent Practice17

(“Extensions” will not be tested on, but are helpful to deepen our understanding and make excellent bonus questions.)Station 1 – Electron Configuration & The Rules Write the full ground state electron configurations and orbital notations (showing sublevels in order of increasing energy) for the following:

# of e- Element (atom) e- configuration Orbital Notations/ diagrams

1._____lithium ________________________________

2._____calcium _______________________________

3._____nitrogen ______________________________

4._____chlorine ______________________________

Write the abbreviated ground state electron configurations for the following:

# of electrons Element

5.______nitrogen ________________________________________

6.______chlorine ________________________________________

7.______iron ________________________________________

8. ______aluminum _______________________________________

Write a ground state electron configuration for these ions.

9.O2-: ______________________________________________________

10.K+: ________________________________________________________

11.Co3+: ______________________________________________________

For the following electron configurations determine the possible elements (or possible ions) they may represent.

12.1s22s22p63s23p64s23d104p4 _________________________________

13. [Kr] 5s24d105p3 __________________________________________

Determine the errors in the configurations. Then, identify which rule has been violated.

14.1s22s22p63s23p64s24d104p5 ____________________ Rule violated:

15.1s22s22p63s33d5____________________ Rule violated:

16.[Ra] 7s25f8____________________ Rule violated:

17.[Kr] 5s24d105p5____________________ Rule violated:

Define the principles and rules.

18. Aufbau principle

19. Pauli-exclusion principle

20. Hund’s Rule

Check your answers and try the mini quiz at the station. Write your question.

Mini quiz answers:

1. _______________________________________________________

2. _______________________________________________________Your question & the correct answer:

Station 2 – Anomalies and Excited States Write the full ground state electron configurations and orbital notations (showing sublevels in order of increasing energy) for the following anomalies:

# of e- Element (atom) e- configuration Orbital Notations/ diagrams

1._____copper ________________________________

2._____chromium _______________________________

Although the IB curriculum will not assess you on the following exceptions, predict the ground state electron configurations for the following metals that are also anomalies to the Aufbau Principle. Write the abbreviated configuration and justify your answer.

# of e- Element (atom) abbreviated e- configuration

3._____ Ag: ______________________________________________________

Why does this occur?

4._____ Mo: ______________________________________________________

Why does this occur?

Excited states of atoms can show electron behavior using electron configurations.

5. How can an electron become “excited”?

Write an excited state full electron configuration for each.

6. Al

7. Ar

8. K

9. C

10. See the example, and determine which elements the excited electron configurations represent.

a. 1s22s22p53s2 atom: ____Na___ what type of “jump” has occurred? __n=2n=3_

b. 1s22s22p64s1 atom: __________ what type of “jump” has occurred? ____________

c. 1s22s22p44s2 atom: __________ what type of “jump” has occurred? ____________


Mini quiz answers:

1. _______________________________________________________

2. _______________________________________________________Your question & the correct answer:

Station 3 – Extension: Trouble with the Aufbau principleRead the article and consider the following questions.

1. Summarize the trouble with the Aufbau principle in 1-3 sentences.

2. Based on what we have studied about the anomalies and the Aufbau principle and combining what you have read, why then do the anomalies occur?

3. Why though, is it still useful to use the Aufbau principle?


Mini quiz answers:

1_______________________________________________________

2_______________________________________________________Your question & the correct answer:

Station 4 –Electron Behavior, Energy, and Light As you read through the review text, answer the questions.

About 300 years ago, Sir Isaac Newton saw a beam of sunlight through a glass prism. He discovered that light is made up of a spectrum of seven distinct visible colors. This spectrum of colors always appears in the same order. You can see this color spectrum (Red, Orange, Yellow, Green, Blue, Indigo, Violet and all the colors in between) when you look through a diffraction grating. There are two color ranges that are not visible to our eyes in this spectrum: below red is infra-red and above violet is ultra-violet.

1. What is the wavelength range for:

a. Infrared ____________________ nm

b. Reds ____________________ nm

c. Oranges ____________________ nm

d. Yellows ____________________ nm

e. Greens____________________ nm

f. Blues____________________ nm

g. Violets ____________________ nm

h. UV____________________ nm

The color of a solid object depends on the colors of light that it reflects. A red object looks red because it reflects red light and absorbs all other colors. A blue object looks blue because it reflects blue light and absorbs all other colors. A white object reflects all colors of light equally and appears white. A black object absorbs all colors and reflects no visible light and appears black. Just like when you color with too many colors in one area with crayons or markers, all colors are absorbed, none are reflected and it appears black!

Type of Spectrum

Photographic example

Continuous (or continuum)

Absorption (dark line)

Emission (bright line)

2. So, emission spectra lines of color show _____________________________ while absorption

spectra lines show____________________________.

Explanation of visible light at the electronic level:

What do fireworks, lasers, and neon signs have in common? In each case, we see the brilliant colors because the atoms and molecules are emitting energy in the form of visible light. The chemistry of an element strongly depends on the arrangement of the electrons. Electrons in an atom are normally found in the lowest energy level called the ground state. However, they can be "excited" to a higher energy level if given the right amount of energy, usually in the form of 3. ______________________ or _______________________. Once the electron is excited to a higher energy level, it quickly loses the energy and "relaxes" back to a more stable, lower energy level called the 4. ____________________ state. If the energy released is the same amount as the energy that makes up visible light, the element produces a color. The visible spectrum, showing the wavelengths corresponding to each color, is shown below:

The important thing to know about absorption and emission lines is that every atom of a particular element will have the same pattern of lines all the time. And the spacing of the lines is the same in both absorption and emission, only emission lines are added to the continuum, while absorption lines are subtracted.

5. In the flame test, you observe a green light from a copper compound, a lilac color from a potassium flame, and a deep red flame from a lithium compound. Rank them in order of increasing energy, and justify your answer.

4. Referring to question 5, which compound emitted the smallest wavelength of light?

5. What is the relationship between energy, frequency, and wavelength?

6. Do you think we can use the flame test to determine the identity of unknowns in a mixture? Why or why not?

7. How are electrons “excited” in a flame test experiment?

8. What particles are responsible for the production of colored light?

9. Why do different chemicals emit different colors of light?


Mini quiz answers:

1_______________________________________________________


Station 5 –Emission Spectra and Analysis Use the spectra at the station to answer the following questions.

1. What do the different colors in a line spectrum represent? Why are the spectra for each element unique?

2. It has been calculated that the observed colors in a hydrogen atom correspond to the relaxation of the electron from a higher energy level to the second energy level. Which color corresponds to 62 transition? 5 2 transition? 42 transition? 32 transition? Explain your reasoning.

3. Which element produced the largest number of lines? Which element produced the smallest number of lines? Explain why elements produce different numbers of spectral lines.

4. What is the converging of spectral lines? Refer to an example to explain.

Using the simplified spectra on the next page, answer the questions.

5.List all elements present in unknown sample W.

6.List all elements present in unknown sample X.

7.List all elements present in unknown sample Y.

8.List all elements present in unknown sample Z.


Mini quiz answers:

1_______________________________________________________


Station 6 –Calculating Photon Energy Max Planck theorized that energy was transferred in

chunks known as quanta, equal to hv. The variable h is a constant equal to 6.63 × 10 -34 J·s and the variable v (or sometimes f) represents the frequency in s-1 or Hz (hertz). This equation allows us to calculate the energy of photons, given their frequency. If the wavelength is given, the energy can be determined by first using the wave equation (c = λv) to find the frequency, then using Planck’s equation to calculate energy.

Use the equations above to answer the following questions.

1. Ultraviolet radiation has a frequency of 6.8 × 1015 s-1. Calculate the energy, in joules, of the

photon.

2. A sodium vapor lamp emits light photons with a wavelength of 5.89 × 10 -7 m. What is the energy

of these photons?

3. One of the electron transitions in a hydrogen atom produces infrared light with a wavelength of

7.464 × 10 -6 m. What amount of energy causes this transition?

4. What is the wavelength and frequency of photons with an energy of 1.4 × 10 -21 J? Where does it fall on the electromagnetic spectrum?

5. A photon emits a bright aqua light with a wavelength of 489 nm. What is the energy of this photon in Joules?


Mini quiz answers:

1_______________________________________________________


Station 7 – The Bohr Model and Spectra Bohr applied Planck’s concept of photon energy to the line spectra of elements. When elements are excited they emit radiation at fixed wavelengths. He proposed that only certain energy levels are allowed within the structure of an atom. Electrons can move between these energy levels. The light emitted by the elements is a measure of the energy gap between the two electronic states.

1. From the image of the hydrogen electron behavior, what color would we see for the electron jump n=1 n=2?

2.From the image of the hydrogen electron behavior, what color would we see for the electron jump n=1 n=3?

3.From the image of the hydrogen electron behavior, what color would we see for the electron jump n=1 n=4?

Extension: Relating Bohr’s model to photon energy. Use the formulas, and try these!

6. Calculate the ΔE for the n = 3 to the n = 1 transition in hydrogen. Where on the EMS would this appear? What does the sign mean?

7. A hydrogen atom in its ground state absorbs light with a wavelength of 102.6 nm. To which excited state would the electron be excited? (Identify nf, the final energy level)


Mini quiz answers:

1_______________________________________________________


Station 8 – Extension: Ionization Energy Predictions Ionization energy is the energy required to completely remove an electron from an atom. This can be thought of as the transition between n = 1 and n = ∞. (∞ representing infinity, meaning the electron has been completely removed)

1. How was our class definition of first ionization energy different from the one provided?

2. Calculate the energy needed to remove the electron from hydrogen in its ground

state.

3. What wavelength of light would be required to remove the electron from

hydrogen in its ground state? Where is this on the EMS?

4.Write an ionization energy equation of hydrogen. (See your notes.)

No mini quiz for this station.

Station 9 – Average Atomic Mass and Isotopes

1. The term “average atomic mass” is a _________________________average, and so is calculated

differently from a “normal” average. Explain how this type of average is calculated.

2. Define an isotope. Be specific.

3. How do the physical properties of isotopes differ? Be aware – number of neutrons is NOT a testable physical property, but it does influence the physical properties.

4. The element copper has naturally occurring isotopes with mass numbers of 63 and 65.

The relative abundance and atomic masses are:

69.2% for mass of 62.93amu

30.8% for mass of 64.93amu.

Calculate the average atomic mass of copper.

5. Calculate the average atomic mass of sulfur if 95.00% of all sulfur atoms have a mass of 31.972amu, 0.76%

has a mass of 32.971amu and 4.22% have a mass of 33.967 amu.

6. Naturally occurring strontium consists of four isotopes, Sr-84, Sr-86, Sr-87 and Sr-88.

Below is the data concerning strontium:

Sr-84Sr-86Sr-87Sr-88

0.56%9.86%7.00% 82.58%

What is the average atomic mass of strontium?

7. Boron exists in two isotopes, boron-10 and boron-11. Based on the atomic mass, which isotope should be more abundant?

8. Iodine is 80% 127I, 17% 126I, and 3% 128I. Calculate the average atomic mass of iodine.

9. Hydrogen is 99% 1H, 0.8% 2H, and 0.2% 3H. Calculate its average atomic mass.


Mini quiz answers:

1_______________________________________________________


Station 10 – Calculating Percentage Abundances (Remember, percentages must be in decimal form during your calculations…)

1. Bromine has two naturally occurring isotopes. Bromine-79 has a mass of 78.918 amu and is 50.69% abundant. Using the atomic mass reported on the periodic table, determine the mass of bromine-81, the other isotope of bromine. Show all work.

2. Antimony has two naturally occurring isotopes. The mass of antimony-121 is 120.904 amu and the mass of antimony-123 is 122.904 amu. Using the average mass from the periodic table, calculate the abundance of each isotope. Show all work.

4.Strontium consists of four isotopes with masses of 84 (abundance 0.50%), 86 (abundance of 9.9%), 87 (abundance of unknown), and 88 (abundance of unknown). Using the average atomic mass of strontium, calculate the percentage abundances of the heaviest two isotopes. Show all work.


Mini quiz answers:

1_______________________________________________________


Review Packet – Independent Practice

Check your answers at http://shakeribchem.weebly.com

#eAtome configurationorbital diagram

1._____ oxygen ________________________________ _________________________________

2._____potassium ____________________________ _________________________________

3._____hydrogen _____________________________ __________________________________

4. _____ neon __________________________________ ________________________________

5. _____ phosphorous ___________________________ _________________________________

Abbreviated configurations

6.______zinc ________________________________________

7.______barium ________________________________________

8. ______bromine ________________________________________

9. ______magnesium _______________________________________

10. ______fluorine __________________________________________

Electron Configuration Elements (atoms) and Ions

Write the electron configuration and orbital notations for the following Atoms and ions:

Element / Ions

Atomic number

# of e-

Electron Configuration/Orbital Diagrams

F1-

Na1+

Al3+

Cl1-

Br1-

Mg2+

E.In the space below, write the full (unabbreviated) electron configurations of the following elements:

1)sodium________________________________________________

2)iron ________________________________________________

3)bromine ________________________________________________

4)barium ________________________________________________

5)neptunium ________________________________________________

F.In the space below, write the Noble Gas (abbreviated) electron configurations of the following elements:

6)cobalt________________________________________________

7)silver________________________________________________

8)tellurium________________________________________________

9)radium________________________________________________

10)lawrencium________________________________________________

11)manganese________________________________________________

12)silver________________________________________________

13)nitrogen________________________________________________

14)sulfur________________________________________________

15)argon________________________________________________

G. Given an element with atomic number 11, provide the following information:

a. How many electrons will fill each of the following shells and list their subshells.

1st shell:

2nd shell:

3rd shell:

b. Is this element likely to form a cation or anion?_________

c. How do you know?

d. What charge will the ion formed by this element have?

H. Explain, based on electron configuration, why the noble gases are so unreactive. Use helium and neon as examples to illustrate your explanation.

I. Determine which of the following electron configurations are not valid, and why.

a. 1s22s22p63s23p64s24d104p5

b. 1s22s22p63s33d5

c. [Ra] 7s25f8

d. [Kr] 5s24d105p5

e. [Xe]

J. Is light a particle or a wave?

Is light composed of waves or of particles? If light is waves, then one can always reduce the amount of light by making the waves weaker, while if light is particles, there is a minimum amount of light you can have - a single ``particle'' of light. In 1905, Einstein found the answer: Light is both! In some situations, it behaves like waves, while in others it behaves like particles.

This may seem odd. How can light act like both a wave and a particle at the same time? Consider a duck-billed platypus. It has some duck-like properties and some beaver-like properties, but it is neither. Similarly, light has some wavelike properties and some particle like properties, but it is neither a pure wave nor a pure particle.

A wave of light has a wavelength, defined as the distance from one crest of the wave to the next, and written using the symbol . The wavelengths of visible light are quite small: between 400 mm and 650 nm, where 1 nm = 10-9 m is a ``nanometer'' - one billionth of a meter. Red light has long wavelengths, while blue light has short wavelengths.

A particle of light, known as a photon, has an energy E. The energy of a single photon of visible light is tiny, barely enough to disturb one atom; we use units of “electron-volts”, abbreviated as eV, to measure the energy of photons. Photons of red light have low energies, while photons of blue light have high energies.

The energy E of a photon is proportional to the wave frequency v

E = h v

where the constant of proportionality h is the Planck's Constant, h = 6.626 x 10-34 J s.

Also, the relationship between frequency and wavelength can be defined as:

v= c

λ

where c is the speed of light (3×108 metres per second). So, photons still have a wavelength. A famous result of quantum mechanics is that the wavelength relates to the energy of the photon. The longer the wavelength, the smaller the energy. For instance, ultraviolet photons have shorter wavelengths than visible photons, and thus more energy. This is why they can give you sunburn, while ordinary light cannot.

One means by which a continuous spectrum can be produced is by thermal emission from a black body. This is particularly relevant in astronomy. Astronomical spectra can be combination of absorption and emission lines on a continuous background spectrum.

One convenient method of exciting atoms of an element is to pass an electric current through a sample of the element in the vapor phase. This is the principle behind the spectrum tubes. A spectrum tube contains a small sample of an element in the vapor phase. An electric discharge through the tube will cause the vapor to glow brightly. The glow is produced when excited electrons emit visible light energy as they return to their original levels.

When visible light energy from a spectrum tube is passed through a diffraction grating, a bright line spectrum, or line-emission spectrum is produced. Each element has its own unique emission spectrum by which it can be identified, analogous to a fingerprint. Such a spectrum consists of a series of bright lines of definite wavelength. Each wavelength can be mathematically related to a definite quantity of energy produced by the movement of an electron from one discrete energy level to another. Thus, emission spectra are experimental proof that electrons exist in definite, distinctive energy levels in an atom.

Refer to this text to answer the following questions.

1. What is the difference between a line spectrum and a continuous spectrum?

2. Each line in the emission spectrum of the hydrogen corresponds to an electromagnetic radiation with a specific wavelength. Match the 4 observed colors with the following wavelengths: 410 nm, 434 nm, 486 nm, and 656 nm.

3. How are electrons “excited”? What happens when the electrons “relax”?

4. Each element has its own unique line emission spectrum, just like fingerprints. Explain how this technique can be used to determine the elemental composition of stars.

5. How can the difference in the brightness of spectral lines be explained?

6. According to the modern theory of the atom, where may an atom’s electrons be found?

7. State the equation used to determine the energy content of a packet of light of specific frequency.

8. What form of energy emission accompanies the return of excited electrons to the ground state?

9.Explain, in terms of electron transition, how bright-line spectra are produced by atoms.

K. Solving for photon energy

1. Find the energy, in joules per photon, of microwave radiation with a frequency of 7.91 × 10 10 s-1.

2. Find the energy in kJ for an x-ray photon with a frequency of 2.4 × 10 18 s-1.

3. A ruby laser produces red light that has a wavelength of 500 nm. Calculate its energy in joules.

4. What is the frequency of UV light that has an energy of 2.39 × 10 -18 J?

5. The frequency of violet light is about 7.495x10 14 Hz. What is the wavelength of this radiation?

6. What is the frequency of a photon that has a wavelength of 1428 nm? What type of radiation is this?

7. A popular radio station broadcasts at 107.9 MHz (M = 10 6 ). Find the wavelength of this radiation, in meters, and the energy of one of these photons, in J. What type of radiation is this?

8. What is the energy of a photon with:

a) a wavelength of 58 nm? What type of radiation is it?

b) a wavelength of 0.065 cm? What type of radiation is it?

9. Which of the following are directly related?

a) energy and wavelength

b) wavelength and frequency

c) frequency and energy

L. Planck recognized that energy is quantized and related the energy of radiation

(emitted or absorbed) to its frequency.

ΔE = n h ν where n = integer and h = Planck's constant = 6.626 x 10 -34 J s

1. What is the energy needed to remove the remaining electron from He+ in its ground state? Is it easier or harder to remove the electron from He+ than from H?

M. Average Atomic Mass

1. Calculate the average atomic mass of bromine. One isotope of bromine has an atomic mass of 78.92amu and

a relative abundance of 50.69%. The other major isotope of bromine has an atomic mass of 80.92amu and a

relative abundance of 49.31%.

2. Lithium-6 is 4% abundant and lithium-7 is 96% abundant. What is the average mass of lithium? Try this WITHOUT a calculator!

3.Here are three isotopes of an element: 12C13C 14C

a.The element is: __________________

b.The number 6 refers to the _________________________

c.The numbers 12, 13, and 14 refer to the ________________________

d.How many protons and neutrons are in the first isotope? _________________

e.How many protons and neutrons are in the second isotope? _________________

f.How many protons and neutrons are in the third isotope? _________________

4.Complete the following chart:

Isotope name

atomic #

mass #

# of protons

# of neutrons

# of electrons

potassium-37

oxygen-17

uranium-235

uranium-238

boron-10

boron-11

5. Argon has three naturally occurring isotopes: argon-36, argon-38, and argon-40. Based on argon’s reported atomic mass, which isotope exist as the most abundant in nature? Explain without calculations.

6. Calculate the atomic mass of lead. The four lead isotopes have atomic masses and relative abundances of 203.973 amu (1.4%), 205.974 amu (24.1%), 206.976 amu (22.1%) and 207.977 amu (52.4%).

7. Titanium has five common isotopes: 46Ti (8.0%), 47Ti (7.8%), 48Ti (73.4%),

49Ti (5.5%), 50Ti (5.3%). What is the average atomic mass of titanium?

8. What is the atomic mass of hafnium if, out of every 100 atoms, 5 have a mass of 176, 19 have a mass of 177, 27 have a mass of 178, 14 have a mass of 179, and 35 have a mass of 180.0?

Page 1 of 24

Documents

Electron Configuration Elements (atoms) and Ions · Web viewArgon has three naturally occurring isotopes: argon-36, argon-38, and argon-40. Based on argon’s reported atomic mass,