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Chemistry > Topic 2: Atomic Structure > Flashcards

Topic 2: Atomic Structure Flashcards

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1
Q

Proton
a) Relative charge
b) Relative mass
c) Location

A

a) +1
b) ~ 1
c) Nucleus

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2
Q

Neutron
a) Relative charge
b) Relative mass
c) Location

A

a) 0
b) ~ 1
c) Nucleus

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3
Q

Electron
a) Relative charge
b) Relative mass
c) Location

A

a) -1
b) 1/1836
c) Outside the nucleus in the electron cloud

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4
Q

Nuclear Symbol (aka Standard Atomic Notion)

A

Shows the chemical symbol, the mass number and the atomic number of the isotope.

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5
Q

Atomic Number (Z).

A
  • # Protons of an atom
  • Same atomic number for all the elements
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6
Q

Mass Number (A).

A

of protons + # of nuetrons

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7
Q

Isotope

A

Atoms with the same number of protons but different number of neutrons.

(They are different versions of the same element. They behave the same chemically but physical properties may vary.)

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8
Q

Relative atomic mass (Ar)

A

Ratio of the average mass of the atom to the unified atomic mass unit

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9
Q

Radioisotopes

A

Unstable form of a chemical element that releases radiation as it breaks down and becomes more stable.

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10
Q

Three types of nuclear radiation

A

a) Alpha
b) Beta
c) Gamma

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11
Q

Alpha radiation

A

Alpha radiation occurs when the nucleus of an atom becomes unstable (the ratio of neutrons to protons is too low) and alpha particles are emitted to restore balance.

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12
Q

Beta radiation

A

Beta radiation (symbol β), is a high-energy, high-speed electron or positron emitted by the radioactive decay of an atomic nucleus during the process of beta decay.

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13
Q

Gamma radiation

A

Gamma rays are a form of electromagnetic radiation (EMR).

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14
Q

Uses of radioisotopes

A
  • Nuclear medicine for diagnostics, treatment and research
  • “Chemical clocks” in geological and archaeological dating
  • “Tracers” in biochemical and pharmaceutical research
  • PET (position emission tomography) scans to give 3-D images of tracer concentration in the body and can be used to detect cancers
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15
Q

Isotopic Abundance

A

The percentage of an isotope in a sample of an element.

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16
Q

Mass Spectrometer

A

Used to identify isotopes and their relative abundance.

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17
Q

Stages of mass spectrometer

A
  1. Vaporization
  2. Ionization
  3. Acceleration
  4. Deflection
  5. Detection
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18
Q

Mass spectrometer: Vaporization

A
  • all particles passing through get converted to gaseous state
  • high vacuum here so particles don’t collide with air
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19
Q

Mass spectrometer: Ionization

A
  • the gaseous atoms are bombarded with high-energy electrons
  • to generate positively-charged species
    e.g. X (g) + e- -> M+ (g) + 2e-
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20
Q

Mass Spectrometer: Acceleration

A
  • the ions are attracted to positively-charged plates
  • thus they’re accelerated in the electric field
  • so they all have the same KE
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21
Q

Mass Spectrometer: Deflection

A
  • the positive ions are deflected by an electromagnetic field
  • degree of deflection depends on mass-to-charge ratio
  • high deflection: low mass, high charge
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22
Q

Mass Spectrometer: Detection

A
  • the beam of ions passing through the detector plate is electrically detected
  • species of a particular m:z ratio are identified
  • results are called “mass spectrum”
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23
Q

Electromagnetic Spectrum

A

A spectrum of wavelengths comprised of the types of electromagnetic radiation

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24
Q

Properties of electromagnetic radiation

A
  • has electric and magnetic fields that oscillate perpendicularly to each other and to the direction of travel
  • behaves like both a particle and like a wave
  • velocity of EM waves = velocity of light
  • can travel in a vacuum
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25
Q

Characteristics of red light

A
  • High wavelength
  • low frequency
  • low energy
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26
Q

Characteristics of violet/purple light

A
  • low wavelength
  • high frequency
  • high energy
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27
Q

Electromagnetic Spectrum (EMS)

A

Is a spectrum of wavelengths that comprimise the various types of electromagnetic radiation.
Eg. Visiable light, radio waves, X-rays, microwaves, infrared radiation (IR), ultraviolet radiation (UV) etc.

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28
Q

Relation of energy, wavelength, and frequency

A

a) Energy is inversely proportional to wavelength
b) Energy is proportional to frequency
c) Energy, E, of radiation is inversly proportional to the wavelength (Formula 🠪 E ∝1/vλ)
d) Wavelength (λ) is inversly related to frquency (v). (Formula 🠪 c=vλ, c = speed of light = 3.00 x 10^-8 m s^-1)

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29
Q

Two types of emission spectra that can be produced by matter

A
  1. A continuous spectrum which contains every wavelength in a particular region of the electromagnetic spectrum. (Eg. White light passing through a prism)
  2. A line emission spectrum indicates only particular wavelengths as discrete lines.
30
Q

White light

A

As perceived by the human eye consists of many colors or wavelengths of light.

31
Q

Two types of line spectra

A
  1. Absorption (bright-line emission) spectrum: A series of bright lines of light produced or emitted by a gas excited by heat or electricity (High v)
  2. Emission (dark-line emission) spectrum: A series of dark lines on a continuous spectrum; produced by placing a cool gas between the continuous spectrum source (white light) and the observer and noting which wavelengths get absorbed by the gas sample.
32
Q

Line spectrum and element

A

Each element has its own characteristic line spectrum which can be used to identify the element

33
Q

Quantization of energy

A
  • Electromagnetic radiation comes in discrete parcels
  • One pocker of enegry or photon is released for each electron transmission
  • The greater the energy of the photon the smaller the wavelength
  • This is shown by E hv = hc/λ: or E is inversely proportional to λ. E is energy of the photon, h is Plank’s constant, v is frequency of the radiation.
34
Q

Hydrogen line spectrum

A

Discrete lines which converge at higher energies and form a continuum

35
Q

Ionization in line spectrum

A

Beyond the convergence limit the electron can have any energy and is not longer in the atom.

36
Q

Series of lines in the hydrogen line emission spectrum

A

a) Balmer series
b) Lyman series
c) Paschen series

37
Q

Balmer series

A

a) Visible region
b) Electronic transitions from upper energy levels back down to the n = 2 energy level.

38
Q

Lyman series

A

a) UV region
b) Electronic transitions from upper energy levels back down to the n = 1 energy level.

39
Q

Paschen series

A

a) Infrared Radiation (IR) region
b) Electronic transitions from upper energy levels back down to the n = 3 energy level.

40
Q

Wavelengths of visible light range in the continuous spectrum

A

400 to 700 nm

41
Q

Emission spectrum

A

A spectrum of the electromagnetic radiation emitted by a source.

42
Q

How is an emission spectrum formed?

A

When light passes through a gas or cloud at a lower temperature than the light source, the gas absorbs at its identifying wavelengths, and a dark-line, or absorption, spectrum will be formed.

43
Q

The Bohr Model of the Atom

A

Bohr’s atomic model pictures electrons in orbit around a central nucleus.
1. Electrons can orbit without losing energy, they are “quantified”
2. The further away the greater the energy of the electron
3. Orbits coverage at higher energies because the energy levels are closer together, forming a continuum and the electron becomes a free electron

44
Q

Atomic Theory

A
  • The lowest energy state for electrons in an atom is its “ground state” (no enrgy is emitted)
  • An atom has a specific, allowable energy level which correspond to the atom’s electrons occypying fixed, circular “orbits” around the nucleus.
  • An atom can gain or lose a specific quantity of energy; when absorbing energy the electron moves to an “excited” state and temporarily occupies a higher energy level
  • Assumes each energy level can hold a maximum number of electrons = 2n^2
45
Q

Successes of the Bohr Atomic Model

A
  • fixed limitation of Rutherford’s model (e-accelerating & losing energy)
  • formed the basis for writing electron arrangements (e.g. P: 2, 8, 5)
  • based on the fundamental idea of quantization with electrons existing in energy levels and included the idea that electrons move from one energy level to another
46
Q

Limitations of the Bohr Atomic Model

A

it successfully explained only one-electron systems. That is, it worked dine for hydrogen and for ions with only one electron, but was unable to explain the emission spectra produced by atoms with two or more electrons.

47
Q

Ionization energy

A

The energy needed to remove an electrin from ground state of a gaseous atom. These energies can also be used to support the model of the atom.

48
Q

Quantum Mechanics

A

Studies motion at the atomic level where the laws of classical physics do not apply since particles behave like waves.

49
Q

Describe the electron according Louis deBroglie

A

Originate the idea that the electron is a particle that displays wave properties.

50
Q

Heisenberg Uncertainty Principle

A

States that it is impossible to determine accurately both the momentum and the position of a particle

51
Q

Evidence that energy levels were further divided into sub-levels

A

Large spaces between individual colors in emission spectra suggest that there are individual energy levels with significant energy differences between them

52
Q

Sub-level

A

A 3-D region for electrons to occupy with subtle energy difference within the main energy level

53
Q

Schrodinger’s Equation

A

a) Developed a mathematical model to describe the atom
b) WAVE EQUATION = mathematical equation integrates the particle-wave nature of the electron.
c) There are many solutions to this equation, and each of these solutions represent particular wave functions, which shows the possible energy states an electron occupy
d) Mathematically, it describes the probability of finding an electron in a region of space at a specific distance from the nucleus. This termed the ELECTRON PROBABILITY DENSITY which can be simply regarded as an electron cloud but can be better described (visualized) as an atomic orbital

54
Q

Atomic orbital

A

Is a region in space where there is a high probability of finding an electron

55
Q

Characteristics of Orbitals

A

a) It is a 3-D region around the nucleus where there is a high probability of finding an electron
b) If an electron has a definable energy, then it can be localized in an orbital of certain size and shape

56
Q

Types of atomic orbitals

A

s, p, d, and f.

57
Q

s atomic orbital

A
  • spherically symmetrical
  • the sphere represents a boundary surface, meaning that within the sphere there is a 99% chance or probability of finding an electron
58
Q

p atomic orbital

A
  • dumbell shaped
  • there are three p atomic orbitals, px, py, pz, all with boundary surfaces conveying probable electron density pointng in differnt directions along the three respective Cartesian axes, x, y, and z.
59
Q

Quantum Mechanical Model of the Atom

A

a) The overall shape of an atom is a combination of all its orbitals (spherical)
b) these orbitals are 3-D and develop (vs. distinct levels in the Bohr model)
c) Each orbital holds 2e- max
d) Electrons can move amongst orbitals if it absorbs (emits) sufficient quanta of energy

60
Q

Aufbau principle

A

The Aufbau principle states that electrons fill the lowest-energy orbital that is available first.

61
Q

Principle Quantum Number (n)

A
  • describes the main enrgy of an electron and the atomic orbital which it occupies
  • the vaue for the energy level expressed as n = {1,2, 3 …. ∞}
62
Q

Orbital diagram

A

Is used to represent the electrons in these atomic orbitals (represent electron configurations).

63
Q

Reasons for removing electron from 4s before 3d levels of 3d-block

A

Since the 4s orbitals has the lower energy, it gets filled first. When 3d orbitals are filled, 4s is no longer lower in energy. Hence, electrons are lost from 4s orbital first, because electrons lost first will come from the highest energy level (furthest away from the nucleus).

64
Q

Spin magnetic quantum number (ms)

A
  • describes the spatial orientation of the spin of the electron
  • In an orbital, the two spin states of the electron are represented by values of + 1/2 and - 1/2
65
Q

Secondary Quantum Number (l)

A
  • describes the sub levels (or sub shells) within a main energy level (shell)
  • the value for the sub shell is expressed as l = {0, 1, 2, …. n-1}
  • (each l value corresponds to a letter designation and name
66
Q

Magnetic Quantum Number (ml)

A
  • refers to the 3-D orientation of the subshells relative to the x-y-z-coordinate system; more simply, these are called orbitals
  • calculated to include whole number values {-l, …, 0,… +l}
67
Q

Pauli Exclusion Principle

A

States that any orbital can hold a maximum of two electrons, and these electrons have opposite spin.

68
Q

Hund’s rule of maximum multiplicity

A

States that when filling degenerate orbitals (orbitals of equal energy), electrons fill all the orbitals singly before occupying them in pairs.

69
Q

Valence Electrons

A

Any electron that exits in an energy level with the highest principle quantum number.

70
Q

Reason for exceptions in Cr and Cu in electron configuration.

A

They are exceptions because it is easier for them to remove a 4s electron and bring it to the 3d subshell, which will give them a half filled or completely filled subshell, creating more stability.

71
Q

Ferromagnetism

A

The basic mechanism by which certain materials (iron &cobalt & nickel) can form permanent magnets.
- Due to the magnetic moment of spinning electrons
- due to many unpaired d-orbital electrons in smaller (period 3) atoms
- due to alignment of “domains”

72
Q

Paramagnetism

A

Weaker magnetic fields due to fewer unpaired e- in atom

Chemistry (12 decks)

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