Unit Three: Atomic Structure

Question 1

Law of definite proportions with example.

Answer 1

A chemical compound always contains exactly the same proportion of elements by mass. Example: oxygen makes up about 8/9 of the mass of any sample of pure water, while hydrogen makes up the remaining 1/9 of the mass.

Question 2

Law of multiple proportion with example.

Answer 2

If two elements form more than one compound between them, then the ratios of the masses of the second element which combine with a fixed mass of the first element will be ratios of small whole numbers. Example: CO and CO2. A fixed mass of carbon, say 100 grams, may react with 133 grams of oxygen to produce one oxide, or with 266 grams of oxygen to produce the other. The ratio of the masses of oxygen that can react with 100 grams of carbon is 266:133 ≈ 2:1, a ratio of small whole numbers. Dalton interpreted this result in his atomic theory by proposing (correctly in this case) that the two oxides have one and two oxygen atoms respectively for each carbon atom. In modern notation the first is CO (carbon monoxide) and the second is CO2 (carbon dioxide).

Question 3

How does the Modern Atomic Theory differ from Dalton’s atomic theory?

Answer 3

1. Dalton’s theory states that atoms are indivisible. Modern atomic theory: atoms are composed of smaller elementary particles.
2. Dalton’s theory states that atoms are indestructible: we now know that atoms can change when they participate in nuclear reactions.
3. Dalton’s theory: an element has only one type of atom. Modern atomic theory: elements can have atoms of more than one type – differring in atomic mass but having the same chemical properties.

Question 4

Who discovered electrons? By which apparatus? What is the charge of this particle?

Answer 4

1. J. J. Thomson (1897) 2. Cathode ray tube. 3. Electron charge: -1, 4. Statistics: Fermionic

Question 5

Rutherford’s gold foil experiment: he bombarded a thin gold foil with ____ particles, which are ______ charged.

Answer 5

Rutherford’s gold foil experiment: he bombarded a thin gold foil with alpha particles, which are positively charged.

Question 6

What happened to the alpha particles in Rutherford’s experiment?

Answer 6

Most of them passed through the gold foil but a few of them rebounded. Around 1 in 8000 alpha particles were deflected by very large angles (over 90°), while the rest passed straight through with little or no deflection. Rutherford concluded that the majority of the mass was concentrated in a minute, positively charged region (the nucleus/ central charge) surrounded by electrons. When a (positive) alpha particle approached sufficiently close to the nucleus, it was repelled strongly enough to rebound at high angles

Question 7

Compare Rutherford’s model of the atom to Thomson’s.

Answer 7

Thomson model: analogous to a plum pudding with the negative charges (the plums) distributed throughout a positive sphere (the pudding).

     Rutherford's model:  

1. the majority of the mass of an atom was concentrated in a minute, positively charged region (the nucleus/ central charge) surrounded by electrons.
2. Atom as mostly empty space: If the nucleus was the size of a marble, the size of the atom would be the size of a football field.
3. That electrons orbited the nucleus, like planets around the sun.

Question 8

Rutherford’s model of the atom was deficient in what respect?

Answer 8

Rutherford suggested that electrons orbit the nucleus; he could not provide an explanation for what kept electrons in orbit around the nucleus.

Question 9

Define electromagnetic radiation

Answer 9

Electromagnetic radiation (EM radiation or EMR) is a form of energy which has a dual wave-particle character. Energy is transmitted by photons which exhibit wave-particle duality as they travel through space.

Question 10

How are λ and ν related in the wave equation?

Answer 10

λν = c where λ is wavelength and ν is frequency and c is the speed of light. Since the speed of light is constant, wavelength and frequency are inversely proportional to each other.

Question 11

Photoelectric effect

Answer 11

The emission of electrons from a metal when light shines on it.

Question 12

What was puzzling about the photoelectric effect?

Answer 12

The wave theory of light predicted that light of any wavelength would be sufficient for the photoelectric effect. However, scientists found that for a given metal no electrons were emitted if the frequency of light was below a certain minimum: no matter how much energy was being delivered to the metal.

Question 13

Draw a diagram illustrating the photoelectric effect.

Answer 13

1. From NYU: Show: Electron
2. Light
3. Photo cathode
4. Anode (+)
5. Meter

Question 14

Define quantum

Answer 14

Quantum: the minimum quantity of energy that can be lost or gained by an atom.

Question 15

Define photon.

Answer 15

Photo: A particle of electromagnetic radiation having zero mass and carrying a quantum of energy.

Question 16

When studying the hydrogen atom emission spectrum, scientists found that it was a line-emission spectrum; not a continous one. How did they explain this phenomenon?

Answer 16

1. Scientists devised the quantum theory to explain this phenomenon.
2. Quantum theory implied that the hydrogen atoms could only exist at fixed, discrete energy states (ground state and then various excited states, analogus to the rungs of a ladder).
3. A hydrogen atom will emit a photon when it transitions from an excited state to its ground state or to a lower energy state.
4. The energy of this photon (E_photon = hν ) is equal to the difference between these two states.
5. If these energy states can only occupy fixed values, emissions will also occur only at specific frequencies (following from the above equation.)

Question 17

Bohr’s model of the atom was based on the Hydrogen Line Spectrum. He proposed that electrons reside in specific energy levels that orbit the nucleus. How are these energy levels analogous to the rungs of a ladder?

Answer 17

1. The amount of potential energy that we can have by standing on a ladder depends upon whether we are on which rung we are on (first, second or ninth – for example).
2. It cannot correspond to an intermediate state because we cannot stand in mid-air.
3. In the same fashion, an electrons can occupy one orbit or another – which are analogized to rungs– but not in between.

Question 18

How does the behaviour of electrons resemble the behaviour of waves?

Answer 18

1. Electrons can be considered to be waves confined to the the space around atoms (the de Broglie hypothesis).
2. **Diffraction: **Electrons exhibited diffraction; a phenomenon which also exists in waves (bending when it passes by the edge of an object).
3. **Interference: **Electrons also exhibit interference: the reduction and increase in energy in different areas when waves overlap. Example: the passage of electrons through a crystal.

Question 19

Heisenberg uncertainity principle.

Answer 19

It is impossible to determine both the position and velocity of anelectron (or any other particle) at the same time.

Question 20

What is the quantum theory?

Answer 20

1. Quantum mechanics deals with phenomenon where dimensions are in the order of Planck constant; the realm of atomic and subatomic length scales.
2. Provides a mathematical description of the dual wave-particle nature of energy and matter.
3. A framework which explains the behavior of atoms during chemical bonding.

Question 21

What is the quantum theoretical conception of an electron orbital?

Answer 21

An orbital is a 3-dimensional region around the nucleus that contains an electron’s probable position.

Question 22

Who coined the term atom?

Answer 22

Democritus

Question 23

What are the elements of Dalton’s Atomic Theory?

Answer 23

1. All matter is composed of extremely small particles called atoms.
2. Atoms of a given element are identical in size, mass and other properties.
3. Atoms cannot be subdivided, created or destroyed.
4. Atoms of different elements can combine in simple whole number ratios to form chemical compounds.
5. In chemical reactions, atoms are combined, separated or rearranged.

Question 24

Examples of the Law of Multiple proportions.

Answer 24

1. CO, CO₂
2. NO, N₂O, NO₂

Question 25

Law of conservation of mass.

Answer 25

The atom inventory in a balanced equation should be equal on both sides.

Question 26

Dalton analogized atoms to be like?

Answer 26

Marbles or pool balls. (Reference?)

Question 27

In what ways are Dalton's laws no longer considered to be entirely correct?

Answer 27

1. Subdividable into elementary particles. 2. Two different Isotopes of an element will have different atomic masses. To be sure, the conviction that atoms cannot be subdivided, created, or destroyed into smaller particles when they are combined, separated, or rearranged in chemical reactions is inconsistent with the existence of nuclear fusion and nuclear fission, but such processes are nuclear reactions and not chemical reactions. In addition, the idea that all atoms of a given element are identical in their physical and chemical properties is not precisely true, as we now know that different isotopes of an element have slightly varying weights. However, Dalton had created a theory of immense power and importance​. Wikipedia.

Question 28

Who discovered the charge and mass of the electron?

Answer 28

Milikan (1909).

Question 29

Draw a diagram illustrating Milikan's Oil Drop experiment.

Answer 29

 From Wikipedia.

Question 30

Alpha particle: notation?

Answer 30

α, α²⁺, ₂ ⁴He²⁺

Question 31

Draw a diagram of Rutherford's experiment.

Answer 31

 Excellent explanation at: [Socratic.org](http://socratic.org/questions/how-did-rutherford-s-gold-foil-experiment-disprove-the-plum-pudding-model). Wikipedia: The Geiger–Marsden experiments (also called the Rutherford gold foil experiment) .

Question 32

What were the conclusions of Rutherford's experiment?

Answer 32

1. An atom is mostly empty space. This explains why most alpha particles passed through the gold atoms. 2. The bending of a few atoms can be explained by a tiny positively charged nucleus. 3. The nucleus is made of positive protons. 4. The negatively charged particles are spread around the nucleus. 5. 

Question 33

Why were Rutherford's analysis of the Geiger-Marsden experiment so significant for chemistry?

Answer 33

1. The atom as mostly empty space explains why chemical reactions are possible. 2. Protons concentrated in the nucleus: Explains the positive charge of a cation;. 3. Explains the mass of an atom. (???) 4. *Flaky kind of answer -- check with your teacher what is an acceptable answer.*

Question 34

Chadwick discovered:

Answer 34

Neutrons in 1932. *He may have discovered other things, but this is test prep.*

Question 35

The number of _____ determines the element.

Answer 35

Protons

Question 36

Most chemical properties of atoms are determined by their _____ because \_\_\_\_\_\_\_.

Answer 36

Most chemical properties of atoms are determined by their electron configuration because chemical reactions are mediated by electron-electron interactions.

Question 37

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Answer 37

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Question 38

Atomic mass unit: define

Answer 38

U**nified atomic mass unit** (symbol: **u**) or **Dalton** (symbol: **Da**): one twelfth of the rest mass of an unbound atom of [carbon-12](http://www.wikipedia.org/wiki/Carbon-12 "Carbon-12") in its nuclear and electronic [ground state](http://www.wikipedia.org/wiki/Ground_state "Ground state"),^[1] and has a value of 1.660538782(83)×10⁻²⁷  [kg](http://www.wikipedia.org/wiki/Kilogram "Kilogram").^[2] 1 Da is *approximately* equal to the mass of 1 proton or 1 neutron.

Question 39

Actual and relative mass of an electron; relative charge; its mass number; location in relation to the nucleus?

Answer 39

1. ​Electron mass: 9.10938291 × 10^-31 kilograms 2. In AMU: 5.486 303 7178×10⁻⁴ amu. 3. Relative charge: -1 4. Mass number: 0 5. External:

Question 40

Actual and relative mass of a proton; relative charge; its mass number; location in relation to the nucleus?

Answer 40

1. Actual mass: 1.672621777(74)×10⁻²⁷ Kg 2. Relative mass: 1.007277 amu 3. Relative charge: +1 4. Mass number: 1 5. Location: nucleus

Question 41

Neutron: Actual and relative mass; relative charge; its mass number; location in relation to the nucleus?

Answer 41

1. 1.67495\*10^-27 kg 2. 1.008665 amu 3. Relative charge: 0 4. Mass number: 1 5. Inside the nucleus.

Question 42

\_\_ % of a carbon atom's mass is in its nucleus.

Answer 42

99.97% of a carbon atom's mass is in its nucleus.

Question 43

A neutral atom is one that has:

Answer 43

A neutral atom is one that has no charge. (Better phrasing than equal number of protons and electrons since the reader does not have to make one more mental step to understand the sentence.)

Question 44

Periodic Table Notation vs Isotopic Notation: differences

Answer 44

Periodic Table Notation Isotopic Notation Atomic number: On top of E Subscript-before Avg Mass U: Below E --- Mass number: --- Superscript-before Net charge: --- Superscript-after

Question 45

Electron/dust particle: similarities and differences.

Answer 45

1. Both electrons and dust particles have mass. 2. An electron has no definite size or position; presumably, this is not true for a dust particle. I suppose what is meant is that you can accurately measure a dust particles maximum diameter; this has no conceptual meaning for an electron.

Question 46

Electron/photon: differences.

Answer 46

1. Mass: electrons have, photons do not. 2. Velocity: photons always travel at *c.* No idea what that means for electrons, although we will learn about average electron [drift velocity](http://en.wikipedia.org/wiki/Drift_velocity) in a conductor. 3. Any other differences?

Question 47

Xenon radius: the van der Waals b parameter has units of L/mol: Xe: a (L² - atm/mol²) =4.19 b (L/mol) =0.0510. What is the radius of the Xenon atom?

Answer 47

1. https://answers.yahoo.com/question/index?qid=20121105171144AArsElp1 mole contains 6.02 \* 10^23 atoms of XE 2. b = 0.0510 L / mole \* [1 mole / 6.02\*10^23 atoms] = 8.47 \* 10^-26 Liters / atom 3. Next we have to get rid of the Liters. 4. 1000 Liter = 1 Cubic meters 5. 8.47 \* 10^-26 L / atom \* [1 m^3 / 1000 L] = 8.47 \* 10^-29 cubic meters. 6. Using the formula for a sphere. 8.27\*10^-29 = 4/3 pi \* r^3 r^3 = 2.0225 \* 10^-29 cu meters r = cube\_root(2.0225 \* 10^-29 m^3] r = 2.72 \* 10 ^ - 10 meters. 7. Converting to Angstrom units. 1 angstrom unit = 1\*10^-10 meters. r = 2.72 \* 10^-10 m [1 A / 1 \* 10^-10 m] * *r = 2.72 A** 8. **Google: 216 picometers.**

Question 48

Atomic radius units

Answer 48

Picometer or Angstrom

Question 49

of protons, neutrons and electrons in ³⁹₁₉K⁺¹

Answer 49

1. Protons: 19 2. Neutrons: 39-19 = 20 3. Electrons: 19 -1 = 18

Question 50

of protons, neutrons and electrons in ¹⁶₈O^-2

Answer 50

1. Protons: 8 2. Neutrons: 8 3. Electrons: 10

Question 51

of protons, neutrons and electrons in ⁴¹₂₀Ca⁺²

Answer 51

1. Protons: 20 2. Neutrons: 21 3. Electrons: 18

Question 52

Nuclide vs Isotope: difference?

Answer 52

Is there one? A set of nuclides with equal proton number (atomic number), i.e., of the same chemical element but different neutron numbers, are called isotopes of the element. From wikipedia: A nuclide (from nucleus) is an atomic species characterized by the specific constitution of its nucleus, i.e., by its number of protons Z, its number of neutrons N, and its nuclear energy state.

Question 53

Name the nuclides of hydrogen

Answer 53

1. Protium: 1 proton, 0 neutrons. 2. Deuterium: 1 proton, 1 neutron. 3. Tritium: 1 proton, 2 neutrons

Question 54

What is the political/historical significance of deuterium?

Answer 54

Wi Amongst other things: During World War II, Nazi Germany was known to be conducting experiments using heavy water as moderator for a nuclear reactor design .... might allow them to produce plutonium for an atomic bomb ... led to the Allied operation called the "Norwegian heavy water sabotage", the purpose of which was to destroy the Vemork deuterium production/enrichment facility in Norway. Wikipedia.

Question 55

Political/historical significance of tritium.

Answer 55

Tritium is an important component in nuclear weapons. It is used to enhance the efficiency and yield of fission bombs and the fission stages of hydrogen bombs in a process known as "boosting" as well as in external neutron initiators for such weapons. Wikipidea.

Question 56

Non-nuclear weapons use of tritium:

Answer 56

In groundwater studies tritium measurements give information on the time of recharge to the system; the tritium content of precipitation being used to estimate the input of tritium to the groundwater system. http://www.scienceworldjournal.org/article/view/5048

Question 57

Technetium-99: use in medicine

Answer 57

1. **taken up by osteoblasts** attempting to heal a skeletal injury, or (in some cases) as a reaction of these cells to a tumor (either primary or metastatic) in the bone. 2. Also used to detect dead cardiomyocytes: no uptake in parts of the heart which have suffered ischemic injury.

Question 58

Average atomic mass: calculation method

Answer 58

The weighted average of the atomic masses of the naturally occurring isotopes of an element. Algorithm: 1. Let the percentages/proportions number of atoms of the isotopes be p1, p2 ... etc. 2. ∑p_i = 1 : (make sure that the percentages have been decimalized). 3. Let the atomic masses of each isotope be m_1, m₂ .. etc. 4. The Average atomic mass is: ∑pi•m_i 

Question 59

Naturally occurring copper consists of 69.17% copper-63 (AMU: 62.929599), and copper-64 (AMU: 64.927793). What is the average atomic mass of copper?

Answer 59

For obvious reasons: 0.6917\*62.92959 + (1-0.6917)\*64.927793 = 63.5456359849

Question 60

????? Boron contains 21.00% 10B and 80.00% 11B. What is its average atomic mass?

Answer 60

0.21\*10 + 0.7900\*11 = 10.79 atomic mass units.

Question 61

Calculate the relative abundance, given C-12 (12.0000 amu), C-13(13.0034 amu); AAM = 12.01 amu

Answer 61

1. Let the respective proportions of C-12 and C-13 be p₁ and p2. 2. p1 + p2 = 1; 3. p2 = 1 - p1 4. p1\*m1 + (1-p1)\*m2 = a (where a is the AMU). 5. p1(m1 -m2) + m2 = a 6. p1 (m1-m2) = a-m2 7. p1 = (a-m2)/(m1-m2) 8. = (12.01-13.0034)/(12 -13.0034 ) 9. **0.99003388479 ** 10. **Carbon is 0.99003388479% C-12,** **0.0099661152 C-13.** 11. ** Sig figs: ?**

Question 62

Algorithm for Isotopic relative abundance in an element from mass spectrum.

Answer 62

1. Obtain the spectographic data. 2. The y-axis represents relative intensivty. 3. The x-axis rpresents m/z where m is mass and z is the charge. 4. Each isotope will produce a seperate sharp spike (z is the same, but m is different). 5. How will you know from the location of the spikes which isotope will have a higher mass? 6. The relative lengths of spikes tell us the relative proportions (I think). 

Question 63

What keeps an atom togetther?

Answer 63

1. Electromagnetic (would prefer electrical attraction between opposite charges) force keeps electrons with the atoms. 2. I think it is the strong nuclear force which keeps the protons together in the nucleus. Immensely powerful.

Question 64

Which force tends to cause neutrons to repel each other?

Answer 64

Neutrons have spin 1/2 and therefore obey the Pauli exclusion principle, meaning two neutrons cannot occupy the same space at the same time. When two neutrons' wavefunctions overlap, they feel a strong repulsive force. See http://en.wikipedia.org/wiki/Exchange\_interaction .

Question 65

Stable nucleii neutron:proton ratio.

Answer 65

1. Brief answer: 1.5 neutrons for every proton 2. Elements that have an atomic number (Z) lower than 20 "prefer" to have the same amount of protons and neutrons. 3. Eg: Carbon: A-Z=N; 12-6= 6 ; 1:1 ratio (because there are 6 protons and 6 neutrons). 4. Example: Oxygen: A-Z= N; 16-8=8 1:1 ratio (because there are 8 protons and 8 neutrons). 5. For elements that have atomic numbers from 20 to 83 the ratio is 1.5:1, the reason for this difference is *because of the repulsive force between protons: the stronger the repulsion force, the more neutrons are needed to stabilize the nuclei*.

Question 66

Symbol for frequency and its units

Answer 66

ν measured in Hertz (Hz).

Question 67

Frequency

Answer 67

The number of waves that pass a given point in space over a specific period of time (typically 1 second).

Question 68

Wavelength of a blue-violet laser.

Answer 68

Green (520 nm), blue-violet (445 nm) and red (635 nm).  Wikipedia

Question 69

Caclulate frequency: wavelength of the light from a green laser is 520 nm. What is its frequency?

Answer 69

1. Frequency = c/wavelength 2. c = 299792458 m/s 3. 520 nm = 520 \* 10 exp(-9) = 5.2 \* 10^(-7) 4. Frequency = 576523957.692\* 10⁷ 5. About 5.77 x 10¹⁵ Hertz (?? order of magnitude?)

Question 70

Wavelength: Symbol/unit

Answer 70

Smbol: λ, units: meters.

Question 71

Planck's quantum theory

Answer 71

Objects emit energy in small, specific amounts called *quanta;* the relationship between a quantum of energy and electromagnetic radiation is stated by: *E=hν*. where *h* is the Planck constant: 6.626 x 10^-34 J•s.

Question 72

When excited electrons transition back to a ground state, they give off energy in the form of:

Answer 72

Photons

Question 73

Draw a diagram of the electromagnetic spectrum :-)

Answer 73

 From Wikipedia. 1. Show visible light band 2. Show X-rays. 3. Show radiowaves 4. Show gamma rays.

Question 74

Why are fireworks so colorful?

Answer 74

Fireworks designers put salts of specific metals into their products; when the fireworks are ignited these salts reach very high temperatures and emit light at characteristic frequencies (at their emission spectrums).

Question 75

Sodium: flame test color

Answer 75

Yellow

Question 76

Barium: flame test color

Answer 76

Green

Question 77

Strontium: flame test color

Answer 77

Red

Question 78

Potassium: flame test color

Answer 78

K [Potassium](http://en.wikipedia.org/wiki/Potassium "Potassium") Lilac [](http://en.wikipedia.org/wiki/File:Flammenf%C3%A4rbungK.png)

Question 79

Copper: flame test color

Answer 79

Wikipedia Cu(I) [Copper](http://en.wikipedia.org/wiki/Copper "Copper")(I) Bluish-green Cu(II) Copper(II) (non-[halide](http://en.wikipedia.org/wiki/Halide "Halide")) Green [](http://en.wikipedia.org/wiki/File:Flametest--Cu.swn.jpg "Flame test on copper sulfate")

Question 80

Lithium: flame test color.

Answer 80

Wikipedia. Li [Lithium](http://en.wikipedia.org/wiki/Lithium "Lithium") Carmine; invisible through green glass [](http://en.wikipedia.org/wiki/File:Flammenf%C3%A4rbungLi.png)

Question 81

Evolution of the Atomic Model: diagram. Hard, but if you can draw it accurately, would be cool.

Answer 81

 From "Gale Group" at ggsrv.com

Question 82

What is the crucial difference between the Bohr model and the Quantum view of the atom?

Answer 82

Bohr conceived of electrons in circular orbits (at different energy levels); the quantum view is that electrons belong to a cloud with certain paths having greater probabilities than others. Ie, probability clouds rather than fixed orbits. Quantum theory states that the position of an electron can never be stated with absolute certainty.

Question 83

A ball of mass 0.2 kg is thrown with a velocity of 10 m/s. What is its wavelength? Modern atomic theory, iun.edu.

Answer 83

Classical physics viewed particles and waves as two very distinct entities. de Broglie hypothesized that particles and waves could actually behave like each other. He suggested a simple equation that would relate the two: Particles have momentum (p), waves have wavelengths (l) and the two are related by the equation **l=h/p** where **h=Planck's constant = 6.634x10-34 Js** ** ** and **p=(mass)x(velocity) = **momentum of the particle. **** Which is an incredibly small number.

Question 84

Evolution of atomic theory: essay components.

Answer 84

1. Democritus: coined the term atom. 2. Dalton: first atomic theory based on experiments 3. Thomson: discovered electrons, cathode ray tube experiment. 4. Milikan: found the charge and mass of an electron; "oil drop" experiment. 5. Rutherford: found nucleus, "gold foil" experiment. 6. Chadwick: neutrons in the nucleus. 7. Bohr: electron orbits, experiment: emission spectra? 8. Quantum theory: probabilistic model of electron orbitals.