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Chemistry IB HL > Topic 2: Atomic Structure > Flashcards

Topic 2: Atomic Structure Flashcards

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1
Q

Dalton’s Atomic Theory

A
  1. all matter is made up of tiny particles called atoms
  2. an element consists of atoms of a single type
  3. compounds are a combination of 2 or more types of atoms
  4. atoms can’t be created/destroyed in a chemical reaction, only rearranged
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2
Q

radioactive isotope

A

an isotope of a chemical element that has an unstable nucleus, which emits certain radiations.

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3
Q

uses of radioactive isotopes

A
  • radiocarbon dating: C-14 exists in a set ratio to C-12 in living organisms, and when it dies, the C-14 isotopes decay, altering the ratio
  • radiotherapy: Co-60 is a powerful gamma emitter used to treat cancer
  • medical tracer: I-131 releases gamma and beta radiation and can be used to detect if the thyroid is functioning correctly + treat thyroid cancer
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4
Q

stages of mass spectrometer

A
  1. vaporisation
  2. ionisation
  3. acceleration
  4. deflection
  5. detection
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5
Q

Mass Spectrometer: Vaporisation

A
  • high vacuum so particles don’t collide with air

- all particles are converted to gaseous state

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6
Q

Mass Spectrometer: Ionisation

A
  • 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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7
Q

Mass Spectrometer: Acceleration

A
  • the ions are attracted to positively-charged plates
  • accelerated in the electric field
  • so they all have the same KE
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8
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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9
Q

the conditions in which the particle has high deflection on the mass spectrometer

A

high charge to low mass ratio

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10
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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11
Q

electromagnetic spectrum

A

a spectrum of wavelengths comprised of the types of electromagnetic radiation

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12
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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13
Q

characteristics of red light

A

highest wavelength, lowest frequency

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14
Q

characteristics of purple light

A

lowest wavelength, highest frequency

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15
Q

trends in electromagnetic spectrum

A

as wavelength increases:

  • quanta energy decreases
  • frequency decreases
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16
Q

procedure for absorption spectrum to be produced

A
  1. pass electromagnetic radiation (e.g. light) through a collection of cold gas
  2. atoms will absorb some radiation at a certain frequency
  3. the spectrometer will compare the transmitted radiation to initial radiation and produce the absorption spectrum
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17
Q

observations in absorption spectrum

A
  • continuous spectrum of colours
  • with vertical black lines at seemingly random intervals
  • lines indicate absence of transmitted radiation
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18
Q

procedure for emission spectrum to be produced

A
  1. heat gas with electric sparks

2. observe the output through a spectrometer

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19
Q

what’s observed in emission spectrum

A

a black background with seemingly random vertical lines of colour (follows colour spectrum placements)

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20
Q

Bohr’s model

A
  • electrons move in orbit around protons
  • the forces of attraction are balanced by the acceleration of electrons at high velocity
  • the closer the electron to the nucleus, the more stable the electron
  • if energy was negligible, electrons could travel to any shell
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21
Q

Bohr’s model’s explanation of absorption and emission

A
  • electrons absorb energy in the form of photons
  • they get excited to a higher energy level if the photon’s energy >= energy difference between initial and final energy levels
  • due to decreased stability in higher energy levels, electrons emit energy to go back to ground state
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22
Q

isoelectronic species

A

elements/ions that have the same electronic config.

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23
Q

Heisenberg’s Uncertainty Principle

A

it isn’t possible to measure the position AND velocity of a microscopic particle with 100% certainty.

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24
Q

reasoning behind Heisenberg’s Uncertainty Principle

A

it’s impossible to locate a microscopic body without disturbing its position/velocity

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25
Q

Schrodinger Model

A
  • uses wave functions to describe electron behaviour

- shape of atomic orbitals depend on energy of electrons

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26
Q

principal quantum number

A

represented by “n”

  • describes the main energy level (aka shell) occupied by the electron
  • the higher the value of n = the farther the electron from the nucleus = the higher the energy associated with the shell
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27
Q

angular momentum quantum number

A 
represented by "l"
- describes the shape of the orbital (aka subshell)
- formula: 0 to (n-1)
n being principal quantum no
- s, p, d, f, etc.
28
Q

magnetic quantum number

A 
represented by (ml, l in subscript)
- describes the orientation of an orbital around the nucleus
- formula: -l to l
l being azimuthal/angular quantum no
e.g. the x or y in 2px or 2py
29
Q

spin quantum number

A

represented by “ms” (m subscript s)

  • describes the spin orientation of the electron
  • value: either spin up or spin down (0.5 or -0.5)
30
Q

wave functions

A
  • principal quantum no. (n)
  • azimuthal no. (l)
  • magnetic no. (m1)
  • spin no. (ms)
31
Q

relative masses of subatomic particles

A

proton : neutron : electron

1 : 1 : 0.0005

32
Q

relationship between EM wave velocity, frequency, and wavelength

A

speed = frequency x wavelength

33
Q

line spectrum

A

a representation of light appearing as a series of discrete coloured lines

34
Q

continuous spectrum

A

a spectrum in which there are no gaps, each region blends directly into the next

35
Q

Pauli’s Exclusion Principle

A

no 2 electrons in an atom can have the same values for all 4 quantum numbers

36
Q

Hund’s rule

A
  • electron pairing will not occur in orbitals in the same subshell
  • until all orbitals are filled with at least one electron
37
Q

factors affecting ionisation energy

A
  • effective nuclear charge
  • size of atom
  • shielding effect
38
Q

ionisation energy

A

the amount of energy required to remove the most loosely bound electron from an isolated gaseous atom of an element

39
Q

successive ionisation energy

A

the energy required to remove subsequent electrons from the atom in the gaseous state

40
Q

effective nuclear charge

A

the net positive charge experienced by electrons

41
Q

periodic trends of ionisation energy

A
  • increases along a period, as nuclear charge increases and atomic size decreases
  • decreases gradually down a group, as while nuclear charge increases, atomic size and shielding effect also increases
42
Q

exceptions to periodic trends of ionisation energy

A

IE of Be > B
IE of N > O

because B and O configurations are half-filled, while Be and N configurations are completely filled

43
Q

atomic mass

A
  • top left in a chemical symbol

- no. of protons + no. of neutrons

44
Q

atomic number

A
  • bottom left in chemical symbol

- no. of protons/electrons

45
Q

relative abundance of an isotope

A
  • the fraction of a single existing element
  • that has a specific atomic mass
  • can be found with mass spectrometry
46
Q

how is atomic mass calculated

A
  • multiplying relative abundance of each isotope of an element
  • by its atomic mass
  • and summing up all the products
  • then finding the average
47
Q

what does a mass spectrometer do

A
  • ionises atoms/molecules
  • with high energy electron beam
  • then deflects the ions through a magnetic field
  • based on mass to charge ratio of ion
  • the mass spectrum of a sample shows its relative abundance on y axis and m:z ratio on x axis
48
Q

law of conservation of mass

A

matter is not created nor destroyed in a closed system

49
Q

law of constant composition

A

a pure compound will always have the same proportion (ratio) of the same elements

50
Q

Plum pudding model of atom

A
  • negatively charged electrons embedded

- in positively charged “pudding”

51
Q

Aufbau principle

A
  • electrons are placed into orbitals

- in order of lowest energy first

52
Q

relationship between wavelength and frequency

A

c = νλ

c: speed of light
ν: frequency
λ: wavelength

53
Q

disproportionation reaction

A
  • type of redox reaction

- in which a substance is both oxidised and reduced in the same reaction

54
Q

Planck equation

A 
E(photon) = hν
E = energy
h = Planck constant
ν = frequency

OR

E(photon) = hc/λ
c = speed of light
λ = wavelength
55
Q

why is glass transparent?

A
  • 2 most common elements, Si and O2: SiO2 (Quartz)
  • when heated to b.pt, becomes amorphous solid when cooled (neither liquid nor solid)
  • visible light consists of photons, which are absorbed by electrons
  • amorphous solids can’t absorb photons so they just pass through
  • this is due to the large gap between electron shells of glass atoms
  • photons don’t have enough energy to pass through shells so they just go in and out of the atom
56
Q

concept of energy quantification

A
  • energy can only be absorbed/released in small, discrete packages (quantum)
  • instead of a continuous flow
57
Q

relationship between photon energy and wavelength

A
  • inversely proportional

- the greater the photon’s energy, the smaller its wavelength

58
Q

what is the wave function?

A
  • series of mathematical functions describing electrons
  • associates possible energy states that electrons can occupy
  • represented by ψ
  • ψ^2 denotes the probability of finding an e- in a region of space at a given radius “r” from the nucleus
59
Q

atomic orbital

A
  • the means of describing an electron’s wave functions
  • denotes a region in the space around the nucleus
  • where the probability of finding the electron is maximum
60
Q

difference between orbit and orbital

A
  • orbit: 2-D concept

- orbital: 3-D regions

61
Q

nature of light

A
  • light behaves like waves: the diffraction of light that occurs when passing through a small slit can only be explained with a wave model
  • light behaves like particles: scattering of electrons that occur when light hits a metal surface can only be explained using particle models
62
Q

patterns between energy levels and sub-levels

A

nth energy level is divided into nth sub-levels

63
Q

identification of elements

A
  • atoms of diff elements give out a distinctive colour of light when electric discharge is passed through its gaseous form
  • when heated, metals produce different-coloured flames depending on the element
63
Q

reason for convergence in an atomic spectrum

A

energy levels inside atoms are closer together at higher energies

63
Q

exchange energy

A
  • the shifting of e-s from 1 orbital to another
  • in the same subshell

NOTE: positions can only be changed in linear order, e.g. electrons from the 4th shell can’t move to 2nd shell but the opposite may occur

64
Q

relationship between energy of photon and the kinetic energy & work function of an electron

A

E(photon) = KE(electron) + E(W.F.)

65
Q

difference between work function and ionisation energy

A
  • I.E.: energy required to remove an electron from an atom
  • work function: energy required to remove an electron from a metal’s surface
  • the electrons in metal are a “sea” of electrons, they don’t belong to any one atom

Chemistry IB HL (11 decks)

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