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MCAT GENERAL CHEMISTRY > Atomic Theory > Flashcards

Flashcards in Atomic Theory Deck (65):
1

Three particles that make up an atom.

proton, neutron, and electron

2

Where are the protons and neutrons located?

nucleus of atom.

3

These have comparable mass in an atom, while electrons are ______ the mass of a proton

1) proton, neutron

2) electron

4

The charge an electron and proton carries.

1.6 x 10^-19 (negative for electrons)

5

_____ and _____ mainly account for atomic mass.

protons and neutrons in the nucleus (electrons are essentially massless when compared to the other two)

6

What do neutrally charged atoms have?

Same number of protons and electrons

7

What do the numbers on the picture represent?

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1) Bottom Number: A Number: Mass Number (neutrons + protons)

2) Top Number: Z Number: (Number of protons), which equals number of electrons

8

Atoms of the same element that are chemically similiar but have different atomic masses. What atomic particle do they differ in?

1) Isotope

2) Neutrons

 

Note: Isotopes react the same way chemically and is only usually distinguished by mass seperation (I.e. spectroscopy)

9

Weighted average of all masses of an the isotopes for that element. Also take abundance into account.

Average Atomic Mass

10

How do you find the average atomic mass?

Multiple the atomic mass of each isotope by its respective percentage, then add all numbers found together.

11

Demonsrated the existence of opposite charges (plum pudding/blueberry model) in an atom as well as that electrons have fixed charge-to-mass ratio.

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Thompson experiment.

12

How did thompson come to his two major conclusions?

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1) Direction of deflection changed when orietation changed (in regards to electrons)

2) Magnitude of defelction was constant when field of strength was held constant but orientation was aligned differently

13

What are the major things that can bend an electrical field?

Positive and Negative charges (and muons, but not rays, i.e. gamma rays, unless another medium is introduced)

14

elementary particle similiar to an electron that traces a different path than electrons because it is 200 times as massive as an electron.

Muon

15

Measures charge-to-mass ratio in a charged particle. Accomplished by sending charged particles into a perpendicular electric field and observing the degrees to which it curves.

mass spectrscopy

16

What are several factors that effect the radius and degree to which a particle curves in mass spectrscopy?

1) Velocity/Momentum (increase can cause a decrease in deflection and increase in radius)

2) Charge and amount of charge (increase in charge results in increased deflection and decrease in radius; decrease in charge results in decrease in deflection and increase in curvature)

3) Mass (Increase causes decrease in deflection and increase in curvature, decrease in mass increases deflection and decreases curvature)

4) Magnetic field (as this decreases deflection increases and curvature decreases)

17

What are three things mass spectrscopy can determine?

1) Isotopic Abundance

2) Molecular Mass

3) Fragmentation

18

How can the mass of a neutron be obtain from isotopes?

Mass difference between isotopes

19

What does the Millikan Oil Drop experiment entail?

Isolation of anionic drops

 

1) Oil can (atomizer) ejects drops of fine mist that collide with the possibilty of electron transfer

2) Small fraction of the drops assume a negative charge

3) E-field is aligned to oppose gravity an suspend negative charged drops

4) Uncharged oul drops fall unaffected

20

When droplets are suspended in the Milikan Experriment, what does the electric for of the E-field equal?

Graviation force

21

What were two conclusions of the Milikan experiment?

1) Charge of electron has a fixed numerical value (elementary charge; charge (magnitude) of proton equal charge of electron, but opposite sign)

2) When protons and neutrons are combined, the charge equals zero

22

What does Rutherfor experiement entail?

1) Alpha particle passes through a tiny pore in lead plate.

2) After passing though pore, stream passes through a sheet of gold foil

3) Particles pass through the foil and strike luminescent screen (glows when struck by alpha particles)

4) Some particles are deflected by gold sheet (leaving shadows within the screen that indicated the presence of tiny spheres blocking the path)

23

What did Rutherford's experiment conclude?

1) Atoms have dense nuclei where most of atomic mass is located (protons + neutrons)

2) Solids are made of atoms with a dense nucleus and vast empty space between nuclei (disproving plum pudding theory)

3) Gold has uniformly spaced atoms in their microscopic composition (mass associated with the atom occupies very little space and is not spread out uniformly through material, like plum pudding)

4) Electron mass is so small, there whereabouts are not discernable 

24

This princliple quantifies the idea that it is not possible simultaneoulsy to identify a particle's position and velocity.

Heisenberg's Uncertainty Principle

(Premise is that you can know position OR velocity, just not at the same time.)

25

Premise that electrons occupy specific circular orbitals about the nucleus, causing electrons to have specific energy levels.

Bohr (Atomic Model)

26

Electrons can exist only in _________ orbits (electron shells).

specified

27

What is true about transitions of electrons as you get closer to the nucleus.

More energy is required for transitions of electrons as it gets closer to the nucleus

28

What happens when evergy is absorbed by an atom?

Electron transitions to a higher energy level (excitation)

29

What happens when energy is emitted in an atom?

Electron transitions to a lower enery level

30

As you get farther and farther away from the nucelus, what happens to the electron energy levels?

They become closer and closer (making it easier to transition from one state to another state)

 

(i.e. going from n=2 to n=3 requires more energy than going from n=3 to n=4)

31

What is required to excite an electron from a lower to higher energy level?

Ultraviolet (UV) to visible range of electromagnetic radiation

32

What is the energy given off as an electron drops back down to a lower energy level?

Light

33

Lower energy levels have less _______ energy and more ________ states.

1) Absolute energy

2) Stable states 

34

What happens when there are small gaps between energy levels (when an electron is excited)?

The less energy is given off, the longer the wavelength of the light thats emitted.

35

What is the relationship between the energy of a photon and its wavelength of ligh?

Inversley proportional

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38

Shells farther from the nucleus have a greater radius, and thus a greater capacity to hold electrons. What is the equation for maximum occupancy of electrons?

Number of electrons in shell = 2(n)2

 

N=quantum number

39

What is the effective nuclear charge and how do you calculate it?

Charge that accounts for attraction of valence e- to nucleus and repulsive forces between valence e- and core e- as well as other valence e-

40

What is true about the pathway for each electron in an element?

They all follow a UNIQUE PATH.

41

You can never know the ______ and ______ of an electron simultaneously at once. (You can only know one or the other at a time)

Velocity and position

42

Two opposite spins in electron spin pairs always produce _______ magnetic fields.

43

What do shells represent? What do orbitals represent?

Shells: Energy level of e-

Orbitals: Regions where e- are most likely to be found

44

Atom or molecule that contains at least one unpaired electron.

Paramagnetic species

45

Atom or molecule that contains no unpaired electrons.

Diamagnetic species

46

As the distance from the nucleus decreases, there is an ________ probability of finding an e-.

Increase

47

Describe structure of S-orbitals.

48

Describe P-orbitals.

1) Barbell distribution about the nucleus

2) 3 different p orbitals (x, y, z)

3) 1 node for each p orbital (node-abscence of e- density)

49

Describe d-orbitals.

1) Double barbell distribution about the nucleus 

2) 5 d orbitals (xy, xz, yz, x2y2, z2)

3) 2 nodal planes per d orbital

50

No two e- can have the same set of quantum numbers (at once).

Pauli's exclusion principle

51

Hund's Rule

1) e- fill lower energy levels before higher energy levels 

2) degenerate set of orbitals: have singly occupied orbitals (before 2nd e- pairs in that orbital)

52

What happens in regards to half-filled and filled d orbitals?

Single e- is elevated from a lower energy level to that was already paired (in s-orbital) to produce even distribution of e- in the d-level.

53

Which metals are typically the result of half-filled and filled d-shells?

Atoms that typically have 4 and 9 VE

54

When do excited state e- configurations occur?

when e- absorb energy and move to a higher level than it normally occupies in the ground state.

55

n

Describes shell (1, 2, 3, etc.)

56

l

Describes orbital (s, p, d, etc.)

0-3 from s-f respectively 

57

Ml

Describes axis (x, y, z)

p orbitals (-1, 0, 1 

58

Ms

Describes Spin

59

What is released when an electron relaxes back to its ground state?

60

Absorption

excitation

61

Emission

relaxation

62

Lifetime of an excited state

Amount of time an e- is in an elevated energy level

63

What is true for a compound with a long excited state lifetime?

Can store energy by maintaining a high population of e- in elevated energy states

64

In molecules, their are vibrational energy levels to consider as well as rotational and electronic energy levels. What is the result of this?

ground states and excited state exist as a band of energy levels, not just a single level (multiple possibilities exist for the energy transition. Rather than single line absorption or emission spectrum being formed, a range is formed)

65

What is true about the atomic spectrum of hydrogen?

1) All hydrogens exhibit the same absorption and emission spectra 

2) Absorption/Emission occurs in distinct increments w/ distinct energy levels

3) Only Balmer series has emits visible light