Chapter 2/ 12 Flashcards
Atomic structure
Atom
Smallest particles that have the properties of an element
- each element contains only one type of atom
Structure of an atom
- atoms contain a positively charged dense nucleus composed of protons and neutrons (nucleons)
- surrounding nucleus, electrons exist in energy levels (principal energy levels) or electron shells
- main energy level/ shell is given an integer number, n
- it can hold a maximum number of electrons, 2n^2
n = 1, closest to the nucleus, and of lowest energy
The further the energy level is from the nucleus, the higher its number (n) and higher its energy
Proton
Relative charge: 1+
Relative mass: 1
Location: nucleus
Neutron
Relative charge: 0
Relative mass: 1
Location: nucleus
Electron
Relative charge: 1-
Relative mass: 1/2000
Location: surrounding the nucleus
Atoms of the same element
Every atom of the same element has the same number of protons in its nucleus
- no. of protons in the nucleus gives the atom its identity
Atomic number (Z)
Number of protons in the nucleus of an atom
- tells you the electron configuration of an atom
Mass number (A)
Number of protons + number of neutrons in the nucleus of an atom
Mass of an atom
- mass of an atom is concentrated in the dense nucleus
- mass of electrons is negligible, they’re not taken into account when calculating the mass of an atom
- protons and neutrons have the same relative mass, hence, mass of a particular atom depends on total no. of protons and neutrons present in nucleus
Number of neutrons
Mass number - atomic number
Neutral atom
Number of protons = number of electrons
- opposite charges of proton and electron cancel out, leaving atom electrically neutral with no overall charge
Isotopes
Atoms of the same element that have the same no. of protons in the nucleus but different no. of neutrons
- same atomic no. but different mass no.
Radioisotopes
Many isotopes are radioactive
- radioactive isotopes = radioisotopes
- eg. iodine-131, cobalt-60
- used in radiotherapy (involves treating diseases eg. cancer with ionising radiation)
Properties of isotopes
- same chemical properties because they have the same no. of electrons (so, they take part in exactly the same chemical reactions)
- slightly different physical properties, enables isotopes of same element to be separated
Mass spectrometer
Used to determine the relative atomic mass of an element from its isotopic composition
Steps of mass spectrometer
- Sample to be analysed is vaporised to form a gas
- It’s bombarded by high-energy electrons, producing positive ions, then accelerated in an electric field (produce ions with a 1+ charge)
- Positive ions are deflected in a magnetic field depending on their mass: charge ratio
- Ions with a higher mass: charge ratio are deflected less in the magnetic field than ions with a lower mass: charge ratio
- Positive ions reach the detector, where they produce a mass spectrum
Mass spectrum
Shows relative abundance on y-axis and mass: charge ratio on x-axis
Relative abundance
Percentage of that isotope that occurs in nature
Isotopic mass
For singly-charged ions (Ne+) the mass: charge ratio values give mass no. of isotope detected (isotopic mass)
RAM equation
RAM = (abundance x mass)+(abundance x mass) / 100
Bohr model of the atom
Proposed by Niels Bohr, 1914
- electrons can only occupy certain energy levels within the atom
- can transition between these energy levels by absorbing or emitting exact amounts of energy
Sub-levels
- contain a fixed no. of orbitals
Orbitals = regions of space where there is a high probability of finding an electron
- each orbital has a defined energy state for a given electronic configuration and chemical environment
- it can hold 2 electrons of opposite spin
Energy levels and sub-shells
- Main energy levels are split into sub-levels, assigned a number and letters s, p, d or f
- number refers to main energy level, and letter refers to atomic orbital
Atomic orbital
Represents a region of space where there is a high probability of finding an electron
Four atomic orbitals: s, p, d or f s = 2 p = 6 d = 10 f = 14
Pauli exclusion principle
Two electrons can only occupy the same atomic orbital if they have opposite spins (one is arrow up, one is arrow down)
- two electrons in the same orbital will have opposite spins (one CW and one anti-CW)
Aufbau principle
Electrons fill atomic orbitals of lowest energy first
An atom is made of energy levels (n = 1, 2, 3 etc.) that are split into sub-levels
- electrons fill these sub-levels according to Aufbau’s principle
NB/ 4s sub-level fills before 3d
Filling atomic orbitals
- 1s sub-level has lowest energy, so is filled first
- With a given main energy level, s orbitals are of lower energy than p orbitals, hence, fill first
- Atomic orbitals within a sub-level are of equal energy (degenerate)
- includes three p orbitals in 2p, 3p and 4p sub-levels and five d orbitals in 3d sub-level - There’s an overlap between 3d and 4s sub-levels
- means that 4s sub-level is of lower energy and fills before 3d sub-level
Electron configurations
Show how electrons are arranged in an atom
- number in front of letter is principal quantum number, n, which gives no. of main energy level
- letter refers to sub-level (s, p, d, f)
- number in superscript gives no. of electrons in sub-level
Exceptions to Aufbau’s principle
- Chromium (Cr)
- Copper (Cu)
Electron configuration for Chromium
1s2 2s2 2p6 3s2 3p6 4s1 3d5
Electron configuration for Copper
1s2 2s2 2p6 3s2 3p6 4s1 3d10
Negative ions
Formed when atoms gain electrons
Positive ions
Formed when atoms lose electrons
Electron configuration of first-row d-block elements
4s sub-level is lower in energy than 3d sub-level and is filled FIRST
When atoms form positive ions, 4s electrons are removed first
Orbital diagrams
Used to represent electrons in atomic orbitals
- boxes represent atomic orbitals
- electrons are represented by arrows
- arrows point in opposite directions to represent opposite spins of electrons
Hund’s rule
Electrons fill orbitals in same sub-level singly, before pairing up
- because atomic orbitals in same sub-level are degenerate
NB/ electrons that singly occupy 2p orbitals have parallel spins (indicated by direction of arrows)
Line spectra
Used to identify unknown elements
Wavelength
Distance between two crests in an oscillating wave
Units: distance
Frequency
Number of waves that pass a point in one second
Units:
Energy of electromagnetic radiation depends on its frequency
- a higher frequency means a higher energy and vice versa
- Units: Joules
Frequency and wavelength
- have an inverse relationship
- as f increases, wavelength decreases and vice versa
Speed of light
3.00 x 10^8
c = wavelength x frequency
Electromagnetic spectrum
Right to left: wavelength decreases and frequency increases
- energy of different types of ER increases
- Radio waves = lowest energy, lowest frequency, long wavelength
- Gamma rays= highest energy, highest frequency, short wavelength
Relationship between energy, wavelength and frequency of electromagnetic radiation
Higher energy = higher frequency = shorter wavelength
Lower energy = lower frequency = longer wavelength
Infrared radiation
Longer wavelength, lower frequency and lower energy than visible light
- feels warm on the skin
Visible light
- encompasses radiation of intermediate frequency
- visible to naked eye
- takes on 7 different colours (ROYGBIV)
Ultraviolet region
- encompasses UV radiation
- has a shorter wavelength, higher frequency and energy than visible light
- type of non-visible radiation emitted by the sun, dangerous for human skin
Continuous spectrum
- when white light passes through a prism, a continuous spectrum is produced
- it’s produced only if the light source contains all possible wavelengths (white light)
- results when the gas pressures are higher; lines are broadened by collisions between the atoms until they are smeared into a continuum
- shows all the wavelengths/ frequencies of visible light
Absorption line spectrum
Some of the wavelengths of light are missing, shown by black lines on the coloured background
Dispersion
Separation of white light into its component colours
Emission line spectrum
Characterised by having coloured lines on a black background
- emission lines correspond to photons of discrete energies that are emitted when excited atomic states in the gas make transitions back to lower-lying levels
Formation of emission line spectrum
- Electrons can only exist at certain energy levels (main energy levels or shells)
- By absorbing or emitting energy, electrons can transition between the energy levels
- If a high voltage is passed through a gas, electrons in gaseous atom become excited and transition to higher energy levels
- As electrons fall back down to lower energy levels, transitions are accompanied by emission of energy
- Results in formation of an emission line spectrum
Formation of absorption line spectrum
Produced when electrons absorb energy and transition from lower to higher energy levels
- occurs when light of all wavelengths passes through a cold, dilute gas
- atoms in the gas absorb at characteristic frequencies
- as the re-emitted light is unlikely to be emitted in the same direction as the absorbed photon, this gives rise to dark lines superimposed on the continuous spectrum
Visible light emission line spectrum
Electrons transition from higher energy levels to second main energy level (n=2)
- energy emitted when electrons make these transitions corresponds to wavelength or frequency of visible light
- lines converge towards high-energy end of spectrum
Closer the lines are, higher the energy
Relationship between light and energy
E = h x v
E= energy (joules) h= Planck's constant (6.64 x 10^-34 J s) v= frequency (1/s)
Hydrogen emission spectrum
Simplest emission spectrum of the elements- looking at transition of just one electron per atom
Formation of the hydrogen emission spectrum
- as a voltage is passed through a sample of hydrogen, electrons are excited to higher energy levels
- extent to which any particular electron is promoted depends on how much energy electron absorbs
- the significant part of emission spectrum is which energy level the electron falls back down to
Hydrogen emission spectrum energy transitions
- electron transitions to n=1 energy level emit energy that corresponds to UV radiation (Lyman)
- electron transitions to n=2 energy level emit energy that corresponds to visible light (Balmer)
- electron transitions to n=3 energy level emit energy that corresponds to infrared radiation (Paschen)
NB/ n=1 is the ground state
Emissions and energy levels
- A given atom only emits certain energies
- Line spectra are at characteristic frequencies for a given element
- There are only certain energy levels available for electrons in the atom
- Electrons can transition from one energy level to another, but not somewhere ‘in between’
What is significant about line spectra?
Provides evidence for the existence of electrons in discrete energy levels, which converge at higher energies
Quantum model
Energy transitions of the electrons in atoms and molecules can only be understood using a quantum model
Electron transitions
- each element has a characteristic spectrum resulting from electron energy transitions- demonstrated as electrons in atom are quantised
- when this occurs, they can only occupy certain specific potential energy levels in relation to nucleus of atom
Quantised energy
Only certain energy values are allowed
Emission spectra of elements
Evidence for organisation of electrons into energy levels within atoms
Convergence limit
Important feature of lines in each series is that they converge at higher frequencies. The highest-energy end of each series of spectral lines is the convergence limit.
Ultraviolet series
- energy of convergence limit of UV series corresponds to energy absorbed when an electron transitions from n=1 energy level to n= infinity energy level
- n= infinity energy level no longer feel the electrostatic attraction from nucleus
- at this point, electron can be considered to have left the atom, resulting in formation of a positive ion
Ionisation energy
Energy required for this transition of electron between energy levels
Equation: E=hv
E= energy in J
h= Planck’s constant, 6.63 x 10^-34 J s
v= frequency, 1/s
First ionisation energy
Energy required to remove one mole of electrons from one mole of gaseous atoms to produce one mole of gaseous 1+ ions
Ionisation energy
Ionisation energies are always positive (endothermic)
- energy must be added to overcome electrostatic attractions between nucleus and valence electrons
Second ionisation energy
First IE: Energy required to remove one mole of electrons from one mole of gaseous atoms to produce one mole of gaseous 1+ ions
If an additional mole of electrons is removed from one mole of gaseous 1+ ions = second IE
Successive ionisation energies
The sequence of ionisation energies (energies needed to remove electrons from gaseous atoms or ions, to produce increasingly positive ions)
- it’s possible to remove electrons from an atom until only the nucleus remains
- give information that shows relations to electron configurations
Trends in ionisation energy
IE increase progressively as electrons are being removed from increasingly positive ions
- results in a stronger electrostatic attraction between the nucleus and remaining electrons
- trends in 1st ionisation energy across period account for the existence of main energy levels and sub-levels in atoms
Plotting ionisation energies
- used to determine which energy level or sub-level electrons are being removed from
Graph:
- plot of log of IE against no. of electrons removed from atom/ion
- IE increases as no. of electrons removed increases
Determining which sub-levels the electron has been removed from
- electrons are removed from outermost to innermost energy levels (gradual increase in IE as each electron is removed)
- a larger increase occurs when moving from outer to inner sub-levels
When are emission spectra produced?
Produced when photons are emitted from atoms as excited electrons return to a lower energy level
Continuous spectrum vs. line spectrum
A continuous spectrum has all wavelengths, but a line spectrum has only selected or certain wavelengths
How is a visible spectrum produced?
the visible spectrum is produced when a light source passes through a refracting prism and the light is bent through an angle
- this depends on the wavelength of light passing through
- if the light wavelength is long, it doesn’t deviate as much as a short wavelength
Hence, any source of light consisting of several different wavelengths may be separated and displayed on a screen or the different wavelengths may be detected electronically and displayed
Electron transitions
This movement of electrons between the shells
- When electron transitions take place, energy emitted can be detected and its wavelength measured
- This provides information about the relative energies of the energy shells.
Electromagnetic spectrum
Red:
- high wavelength
- low frequency- not many waves per unit time
- low energy
- infrared radiation
Violet:
- low wavelength
- high frequency
- high energy- photons are high energy photons
- ultraviolet radiation
Limit of convergence
In an emission spectrum, the lines converge at higher energies
- at the limit of convergence, the lines merge forming a continuum
- Beyond continuum, electrons can have any energy
- frequency of the radiation in the emission spectrum at the limit of convergence can be used to determine IE1
Converting between frequency and wavelength
v = c/ʎ
Discontinuities in first ionisation energy across the period
- In Ca, big jump between IE2 and IE3 because corresponds to removing an electron from fully occupied 3p sublevel
- In Ti, big jump between IE4 and IE5 because of the change in energy level. Ti5+ does not occur naturally
- Ionization energy for Ti increases more gradually than Ca because electrons are being removed from 3d and 4s orbitals, which are much closer in energy than 3p and 4d.
- These big jumps in energy occur when removal of electrons from different energy levels n (1,2,3,4,…)
Deducing the group of an element from its ionisation energy data
A large increase in IE between the first and second IE = element is found in group 1
- group 1 elements have 1 electron in valence energy level, requires least energy to remove
Graphical representation (log (IE) vs. no. of IE)
- electrons are removed from energy level furthest from nucleus
- electrons in the outer energy level require less energy to remove due to weaker electrostatic attraction from the nucleus
- electrons in the inner energy level require the most energy to remove due to strong electrostatic attraction from the nucleus (higher up the x-axis)
