Atomic Theory

Unit 3 Development of the Atomic Theory 1. Where is the mass of the atom concentrated? In the nucleus 2. What is located in the nucleus? Neutrons and protons 3. What is the negative particle that orbits the nucleus? electrons 4. What is the sum of the protons and neutrons called? Mass number 5. With a neutral atom, what two items are equal in number? Protons and

electrons 6. What is the term for any charged particle? ion 7. What is the term for the positively charged ion? cation 8. What is the term for the negatively charged ion? anion 9. What type of ion is formed when electrons are gained? anion 10. What type of ion is formed when electrons are lost? cation 11. How does atomic theory today differ from Dalton’s theory? Today we

know about ions and isotopes 12. Which model of the atom is based on the solution to the Schrodinger

equation? Quantum mechanical model How is this different from the planetary model? Electrons are not in fixed orbitals

Match each term from the experiments of J.J. Thomson with the correct description.

13. anode d a. an electrode with a negative charge 14. cathode a b. a glowing beam between electrodes 15. cathode ray b c. an electrode with a positive charge 16. electron d d. a negatively charged particle 17. The diagram shows electrons moving from left to right in a cathode-ray tube. Draw

an arrow showing how the path of the electrons will be affected by the placement of the negatively and positively charged electrodes.

H. Cannon, C. Clapper and T. Guillot

Klein High School

Atomic Structure

18. Indicate the letter of each sentence that is true about atoms, matter and electric charge.

a. All atoms have an electric charge. b. Electric charges are carried by particles of matter. c. Atoms always lose or gain charges in whole-number multiples of a single

basic unit. d. When a given number of positively charged particles combine with an equal

number of negatively charged particles, an electrically neutral particle is formed.

19. Indicate the letter next to the number of units of positive charge that remain if a hydrogen atom loses an electron.

a. 0 b. 1 c. 2 d. 3 The positively charged subatomic particle that remains when a hydrogen atom loses

) _proton____. does a neutron carry? neutral

all central region

23.

c

ts

found

ns

20.an electron is called a(n

21. What charge 22. Indicate the letter of each sentence that is true about the nuclear theory of atoms

suggested by Rutherford’s experimental results. a. An atom is mostly empty space. b. All the positive charge of an atom is concentrated in a sm

called the nucleus. c. The nucleus is composed of protons and neutrons.

The nucleusd. is large compared with the atom as a whole. e. Nearly all the mass of an atom is in its nucleus.

According to Bohr, electrons cannot reside at __ figure bein the low

a. point A c. point b. point b d. point d

24. According to the quantum mechanical model, point D in the above figure represene fixed position of an(a) th electron

(b) the farthest position from the nucleus that an electron can be(c)A position where an electron probably exists (d p) a osition where an electron cannot exist

25. What does the atomic number of an atom represent? Number of proto

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Atomic Structure

Unit 4 e t o s in the atom El c r n

b. neutral c. positive

. The number of protons in a neutral atom having 18 electrons is_18.

in the isotope P-29? 14

ass number of 80. . It has_35_ protons.

5 electrons. 6 neutrons.

ne isotope of hydrogen has a mass of 3.

n atom with a +1 charge has atomic number 19 and mass number 39.

s.

2. How many electrons does it have? 10

Select the best possible response. 1.The type of charge on the nucleus is

a. negative

2 3. How many neutrons are A neutral atom of bromine has a m4 . It has_35_

. It has_45_ 7. Its nuclear composition is_35 protons and 45 neutrons___ . O8. How many protons does it have? 1 9. How many neutrons does it have? 2 A10. It has_19__ protons. 11. It has 18___electron The anion of oxygen has a -2 charge. 1_____.

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Atomic Structure

An atom has a charge of +1 and has 10 neutrons with a mass number of 21.

n atom has 10 protons, 8 neutrons, and 12 electrons.

6. What is its charge? 2-

8. In forming this ion, the neutral atom gained what? 2 electrons

and 16 electrons.

number of 24 and 11 protons.

e? 11

s number 32.

.002%; S-33, 0.76%; S-34, .22%; and S-36, 0.014%. What is the average amu of sulfur?

14. What is its atomic number? 11 15. How many electrons does it have? 10 A 1 17. What is its mass number? 18 1 An atom has 15 protons, 16 neutrons, 19. What is its mass number? 31 20. What is the overall charge on this ion? 1- A neutral atom has a mass 21. How many protons does it hav 22. How many neutrons does it have? 13 An atom with a -2 charge has atomic number 16 and mas 23. It has_16___ protons. 24. It has_18___ electrons. 25. It has__16__neutrons. Calculate the following: 26. There are 4 isotopes of sulfur, S-32, 95432.1 amu

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Atomic Structure

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27. There are three common isotopes of chromium, Cr-50, 4.345% ; Cr-53, 9.50% abundance; what is the average

mu of chromium?

8. Calculate the wavelength of the yellow light emitted by a sodium lamp if n is 5.10 x 1014 Hz (5.10 x 1014 s-1).

88 nm iation with a wavelength of 5.00 x 10-6 cm? In

hat region of the electromagnetic spectrum is the radiation?

cy of 5.80 x

he wavelength in cm of the radiation from the station?

abundance; Cr-52, 83.789%a50.8 amu 2the frequency of the radiatio529. What frequency is radw6.00 x 1015 Hz 30. What is the energy of a photon of green light with a frequen1014 s-1? 3.84 x 10-19 31. Suppose your favorite AM radio station broadcasts at a frequency of 1150 kHz. What is t26100 cm

Atomic Structure Practice Element/I

number number sotope Atomic Mass Protons Neutrons Electrons

1) Zinc 30 65 30 35 30 2) Aluminum 13 27 13 14 13 3) Calcium 20 40 20 20 20 4) Sulfur 16 32 16 16 16 5) Bromine 35 80 35 45 35 6) Gold 79 197 79 118 79 7) Silver 47 108 47 61 47 8) Platinum 78 195 78 117 78 9) Uranium – 236 2 144 92 92 236 910) Plutonium – 246 94 246 94 152 94 11) Potassium – 39 19 39 19 20 19 12) Mercury – 201 80 201 80 121 80 13) Titanium - 48 22 48 22 26 22 14) Titanium - 46 22 46 22 24 22

Atomic Structure Exercise

I. Use the periodic table to compute the number of electrons. Neutrons and protons in the following:

Atomic Structure

A. Cr 24p, 24e, 28n B 1 , D. Ir 5n E. F. Ne p, 10e, II. to det ine the atomic nu rs of t lementsidentify each of the following elements by name. A. 1 Proton B. 4 Protons B Proto D. E. 20 Protons C 30 Pro Zn III s are in nucleu ach o e follo elemeA. ium C. H Magn 2 12IV tro A. ogen C. 16 D. 85Cu 56

. Cl 17p, 7e, 18n C. Mg 12p, 12e 12n

77p, 77e, 11 Si 14p, 14e, 14n 10 10n Moseley used x rays erm mbe he e

H e C. ns

12 Protons Mg a F. tons

. How many proton the s of e f th wing nts Uranium B. Selen elium D. esium

92 34 . Give the number of neu ns in each of the following isotopes

Titanium-46 B. Nitr -15 S34 29 24 8 18

The Spectra of Elements

Everybody knows that a few drops of soup or milk spilled onto a gas burner will change the blue gas flame into a mixture of colors, predominantly yellow. These colors

ubstances.

t absorbing ergy and

ses

ot just un constantly

light in a us. The

avelength () is the distance between successive peaks of a wave. The number of waves that pass a given point in space per second is the frequency (f). The unit is inverse time, sec -1 and is also

can be used to identify the elements present in the substance dropped into the flame.We will observe the colors produced by several known s The color observed in the flame is the result of atoms of the elemenenergy from the flame then reemitting it. The energy absorbed has an enwavelength in the visible region of the electromagnetic spectrum. What your sendetect as light is actually radiation that is part of a continuum that makes up the electromagnetic spectrum. The electromagnetic spectrum is a continuum of energies ranging from the high-energy gamma and x-rays to low-energy radio and microwaves. Radiation includes any energy emitted in all directions from a single source, nnuclear decay. Light is one form of energy that can be radiated; - the sradiates energy into space as its matter is converted into energy. We see thisvariety of colors, depending on the wavelength of the energy when it reacheswaves of light with the longest wavelength (red) are refracted (bent) to a lesser degree than the shorter waves (indigo and violet) as they pass through the atmosphere. The visible range of the spectrum is from about 400 nm to 700 nm. The w

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Atomic Structure

called hertz (Hz). All light moves through a vacuum at the same rate of speed, about 3 X 108 m/s (the speed of light – c). Frequency and wavelength are inversely related to each other, as the product of the two equals the speed of light: c = f. Since the electrons of an atom can only absorb certain amounts of energy, the wavelength of the energy emitted after excitement of the atom ent. When viewed

of

d

e and has an easily distinguishable inner blue cone.

ame and observe the result.

on,

s is characteristic of the elemctrum of energy emitted by through a spectroscope, the spe a sample of a particular

element can be observed.

1. Use the spectroscope to view each of the gas tubes set up in the back of the lab. Each colored band corresponds to a different amount of energy (a different wavelength) of light emitted by an electron in theelement. Each element has a unique spectrum based on the amount energy emitted by its electrons

2. Record the details of the colored lines formed on the scale of your spectroscope and the wavelength of each line. Using the equations we have discussed in class, solve for the frequency and energy of each of thegases observed. If a substance has multiple lined, calculate frequency anenergy for each line. Express each answer to three significant digits. Express the energy in ergs. In the next part of our investigation, we will use a procedure known as a flame test to become familiar with the colors produced by several common elements and use theseobservations to better understand the affect of energy on atoms and ions.

Procedure: a. Light a Bunsen burner. Make sure that the gas burns with

a clear blue flamb. Obtain samples of ions as provided by your teacher.

These will be metal nitrates, dissolved in a solvent to allow easy distribution of the ion.

c. Using the spray bottles, introduce a small amount of each ion (one squirt) into the fl

d. After you have identified the color produced by each imix two of the substances and see if you can detect the presence of each of the substances or if the colors mingle so that one cannot be distinguished from the other.

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Atomic Structure

e. Obta

in an unknown sample and observe the colors produced. Try to identify the element present in the unknown sample.

CNgS

What is the order of the colors in the spectrum from lowest to ighest

avele

ements and the energy given off xci

Further Analysis In Bohr hhave onl the electronelectron ve off light when it jumped from a higher energy level to a lower energy level. The amount of energy wo rent fo

1) The diagram below shows a sketch of some of the possible orbits of the ogen electron and their corresponding energy values. If we think of the

value ber ine, w

olor Key: a = Yellow K = Violet Ca = Yellow-red Cu = Blue-reen r = Deep red Li = Crimson Ba = Green-yellow

1. h wavelength?

2. What is the order of metallic ions from lowest to highest w ngth?

3. What correlations if any do you see between the electron configuration of the metallic el

y eb ting their electrons

1913, Neils ypothesized that the electron in the hydrogen atom is allowed to y certain amounts of energy. These energy levels would be orbits in which in hydrogen would circle the nucleus of hydrogen. He believed that the would move from one energy level to another and would gi

uld be diffe r jumps between different levels.

hydrenergy s as being on a num l hich orbit has the greatest value?

2) When an elebe E2 – E1; or (-3.4 eV) – (-13.6 eV) which would equal 10.2 eV. What would be the amount of energy in the light when an electron jumps

ctron jumps from E2 to E1 the amount of energy in the light would

from:

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Atomic Structure

a) E3 to E4 .66 eV b) E4 to E2 2.55 eV c) E3 to E2 1.89 eV d) E5 to E2 2.86 eV

e

a. How many different energies of light can be emitted from hydrogen when the electron jumps down to E2 from E3, E4, E5, and E6?

b. How many bands of light did you observe when you viewed the hydrogen tube through the spectrometer?

c. How do you think these two observations are related?

4) The colors corresponding to jumps to the E1 level have higher energies or lower, than those to the E2.

a. As the electron jumped from the E6 all the way to the E1 level, how many different energies would be emitted?

b. Did you see all of these bands?c. Bohr did not find bands corresponding to the jumps to the E3 level. Why

In this depend nd another. The number e lectrons occupy

Arrangem

1. ern

3. on

4.It is not possible to predict both position and momentum at the same time

level of an electron is the region around the nucleus where an electron to be found____

3) Using Bohr’s model, we would assume that the electron would only movbetween certain orbitals or energy levels. Every possible jump corresponds tolight of a different energy.

Explain why or why not.

do you think that was?

Conclusions:

i p patterns activ ty we saw the evidence that energy is emitted in articularing on the distance that an electron travels between one orbital a of electrons in an atom and the number of energy lev ls those e

determine the number of bands of light and the color of the light seen. What if any conclusions can you draw about the relative numbers of electrons in the atoms of ou er t st atoms?

ent of Electrons in Atoms What is the difference between the earlier models of the atom and the mod

quantum mechanical model? No fixed orbitals

2. What is a quanta and who developed the concept of quantum energy? Packet of energy that it takes to move an electron from one energy level to the next; Planck

How many quantum numbers are used to describe the energy state of an electrin atom?

a) 1 (b) 2 (c) 3 d) 4 What is the Heisenberg uncertainty principle?

gy5. The eneris likely

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Atomic Structure

6. In general, the higher the electron is on the energy ladder, the _further_ it is to/from the nucleus?

7. A quantum of energy is the amount of energy required to

on in its present energy level c. move an electron from its present energy level to the next higher one

rue or False: The quantum mechanical model of the atom estimates the

hich name describes the major energy levels of electrons?

d) principal quantum number at can occupy a

umber)?

different p orbitals combination of an s and a p orbitals

a. move and electron from its present energy level to the next lower one b. maintain an electr

8. True or False: the electrons in an atom can exist between energy levels. F

9. T

probability of finding an electron in a certain position. F

10. Wa) atomic orbitals b) quantum mechanical numbers c) quanta

11. What formula represents the maximum number of electrons thprincipal energy level (n = principal quantum na) 2n2 b)n2 c) 2n d) n

12. A spherical electron cloud surrounding an atomic nucleus would best represent _. a. an s orbital b. a p orbital c. a combination of two d. a

13. An energy level of n = 4 can hold electrons. (a) 32 (b) 24 (c) 8 14. An energy level of n = 2 can hold

(d) 6 electrons.

(a) 32 (b) 24 (c) 8 (d) 6 15. An electron for which n = 4 has more than an electron for which n = 2.(a) spin (c) energy (b) stability (d) wave nature

ure?

e properties to which

did the Heisenberg uncertainty principle contribute to the idea that trons occupy "clouds," or "orbitals"?

Behavior of electrons in the atom

1. How did de Broglie conclude that electrons have a wave nat

2. Identify each of the four quantum numbers and ththey refer.

3. How

elec

4. Complete the following table.

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Atomic Structure

Principal quantum number, Number of Sublevels Types of Orbitals n 1 1 S 2 2 S, p 3 3 S, p, d 4 4 S, p, d, f

5. Th ns are arrange way in which electro ed around the nuclei of atoms is called

d to find the electron configurations of atoms with the

__b__

.electron configuration

Match the name of the rule userule itself.

6. aufbau principle a. When electrons occupy orbitals of equal til

all the orbitals contain one electron with energy, one electron enters each orbital un

parallel spins __c__ 7. Pauli exclusion

principle b. Electrons enter orbitals of lowest energy first

_a___

4-11

ic orbital may describe at most two ectrons moving in opposite directions

8. Hund’s rule c. an atomel

9. In the shorthand method for writing an elof the superscripts eq

ectron configurational? Number of electron

, what does the sumatom

u s in the

s.

nds

Write the electron configuration and orbital notation for each of the following atom

10. Nitrogen 1s22s22p3

11. Aluminum 1s22s22p63s 1

23p

12. Argon 1s22s22p63s23p6 13. Which guideline, Hund's rule or the Pauli exclusion principle, is violated in the

following orbital diagrams?

Pauli Hu

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Atomic Structure

14. What is the relationship between the principal quantum num

cber and the electron

onfiguration? Tells the number of energy levels

15. How does the figure above illustrate Hund's rule? One in each of the 2p levels

before two in any

16. How does the figure above illustrate the Pauli exclusion principle? If 2 in torbital they spin in opposite directions

17. True of False: The aufbau principle works for every element in the periodic table. F

18. Filled energy sublevels are more _stable___ than partially filled sublevels.

19. Half-filled sublevels are not as stable as _filled____ levels but are more stathan other configurations.

20. Write the electron configuration of the following atoms: a. carbon 1s22s22p2

c. gallium 1s22s22p63s23p64s23d104p1 d. copper 1s22s22p63s23p64s23d9

21. What is an electron dot structure? Shorthand way to show the valence electrons

r :

he

ble

b. potassium 1s22s22p63s23p64s1

22. Draw the electron dot structure of each of the following atoms. a. argon .. : A .. b. calcium Ca: c. iodine . : I : ..

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Atomic Structure

23. Write the electron configurations for these metals and circle the electrons lost when each metal forms a ca 2 2 6 2

ation. . Mg 1s 2s 2p 3s

tion.

b. A1 1s22s22p63s22p1

c. K 1s22s22p63s23p64s1

Match the noble gas with its electron configura _c__ a. 1s2 24. argon __ _

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a 25. helium b. 1s2 2s2 2p6 26. neon c. 1s2 2s2 2p6 3s2 3p6 _b__

_d__ 27. krypton d. 1s2 2s2 2p6 3s2 3p63d10 4s2 4p6

13. What is the electron co

Pseudo nobel gas 14. Write the electron conf 15. Fill in the electron c

nfiguration called that has 18 electrons in the outer energy level and all of the orbitals filled?

iguration for zinc. 1s22s22p63s23p64s23d10

onfiguration diagram for the copper(I) ion.

Electro

ite the urations, orbital nollowing:

a. S

n Review 1. Wr electronic config tation, Lewis dot structure for the fo

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Electron Configuration

b. C c. P d. Ca e. Zn f. Fe

. How many dots would appear in the Lewis electron dot diagram for an atom whose lectron configuration ended 4S2 3d10 4p3 ? 5

. How many unpaired electrons does the Lewis dot structure of N have? 3

. How many pairs of electrons does the Lewis dot structure of O have? 2

y does copper have 1 valence electron? Half filled orbitals are more stable than thers

Charting Oxidation Number

Complete the following chart. You may wish to use the periodic table in your text.

g. A1 2e 3 4 5. Who

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Electron Configuration

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