WJEC AS Chemistry Unit 1.2 - Basic Ideas About Atoms

Atomic number

• number of protons in the nucleus

• number of protons=number of electrons

Mass number

number of protons and neutrons in the nucleus

Isotopes

atoms with the same number of protons but different numbers of neutrons

Types of radioactive emission

• Alpha

• Beta

• Gamma

Alpha particles

Have a nucleus of 2 protons and 2 neutrons, therefore positively charged.

Alpha emissions

• least penetrating of the three types

• Stopped by a thin sheet of paper

• Strongly ionising because they are large, relatively slow and carry two positive charges

Alpha symbol

Effect of electric field on alpha emissions

Alpha particles are deflected towards the negatively charged plate

Effect of magnetic field on alpha emissions

• Deflected in one direction

Beta particles

Fast moving electrons

Beta emissions

• Streams of high energy electrons

• More penetrating than alpha particles

• Can travel through air and paper

• Stopped by a thin layer of metal such as aluminium

Beta symbol

Effect of electric field on beta emissions

Beta particles are deflected towards the positive plate

Effect of magnetic field on beta emissions

Deflected in opposite direction to Alpha particles

Ionisation in radiation

• when the types of radiation pass through matter, they knock electrons out of atoms, ionising them

• Alpha = strongly ionising, transfer happens rapidly and so are the least penetrating

• Gamma rays = weakly ionising, most penetrating of the radiations

• Ionisation involves a transfer of energy from the radiation passing through the matter to the matter itself

Gamma rays

High energy electromagnetic radiation, and therefore no charge

Gamma emission

• most penetrating of the three radiations

• Can pass through several cm of lead or more than a m of concrete

Gamma symbol

Effect of electric field on gamma emissions

• no effect

• undeflected

Effect of magnetic field on gamma emissions

Unaffected

Effect of alpha emissions on mass number and atomic number

• mass number decreases by 4

• atomic number decreases by 2

• Product is 2 places to the left in the periodic table

Effect of beta emissions on mass number and atomic number

• mass number is unchanged

• Atomic number increases by 1

• Product is one place to the right in the periodic table

Effect of gamma radiation on mass number and atomic number

No effect as it’s a form of energy and is not an atomic particle

Alpha radiation equation example

Beta- decay equation example

Electron capture

• where one of the electrons is captured by a proton, turning it into a neutron

• An electron neutrino is emitted (V_e)

• Atomic number decreases by 1 due to the changing of the proton into a neutron

• Electron is placed in reactants rather than products

Electron capture example equation

Positrons

• produced by the splitting of a proton into a neutron and a positron

• Type of B+ particle, this results in a decrease in atomic number

• Reverse of beta emissions

Positron emission (B+ decay)

Atomic number decreases by 1

Positron emissions example equation

Half life

Time taken for half the atoms in a radioisotope to decay or the time taken for the radioactivity of a radioisotope to fall to half its initial value

Number of half life periods equations

Periods = time/half-life

Consequences of radiation on living cells

• ionising radiation can damage the DNA of a cell;

  • damage to DNA may lead to changes in the way the cell functions, which can cause mutations and the formation of cancerous cells at lower doses or cell death at higher doses

• outside the body, gamma radiation is the most hazardous

• inside the body, if alpha particle emitting isotopes are ingested, they are far more dangerous than an equivalent of beta emitting or gamma emitting isotopes

Uses of radioisotopes in health and medicine

Uses of radio-isotopes in radio-dating

Uses of radioisotopes in industry and analysis

Radioactivity

• caused by nuclear instability as a result of an unequal p:n ratio, causing the nuclei to change = radioactive decay

• Optimum p:n = 1:1

• Decay —> stable

Atomic orbital

A region in an atom that can hold up to two electrons with opposite spins

Fixed energy levels/quantum shells

• electrons within atoms occupy fixed energy levels or quantum shells

• Shells are numbered

• The numbers are known as the principal quantum numbers, n

• Lower value of n, closer the shell to the nucleus + lower the energy level

S orbital

• spherical orbital

• Represented by a sphere/circle

• Holds up to 2 electrons

• The 2 electrons have to be of opposite spin

P orbital

• Dumbbell shaped

• Hold up to 6 electrons

D orbital

• can hold up to 10 electrons

• Complex 3D shape

Electronic configuration

The arrangement of electrons in an atom

Filling shells and orbitals with electrons

• electrons fill atomic orbitals in order of increasing energy

• a maximum of 2 electrons can occupy any orbital each with opposite spins

• The orbitals will first fill with one electron each with parallel spins, before a second electron is added with the paired spin

4s and 3d sub shells

• 4s is filled before 3d

• Due to increasingly complex influences of nuclear attractions and electron repulsions upon individual electrons

• The s orbitals have slightly lower energy than the d orbitals of the shells that are below them

Exceptions to the electron configuration rules

Config of;

• chromium

• Copper

Configuration of chromium

Configuration of copper

Ionisation

• The process of removing electrons from an atom

• The energy needed to remove each successive electron from an atom is called the first, second, third, second, etc ionisation energy

First standard molar ionisation energy

The energy required to remove one mole of electrons from one mole of gaseous atoms to produce one mole of gaseous ions. Under standard conditions, 25 degrees C and 1 atm

Equation of first ionisation energy

Factors affecting ionisation energy

• The size of the positive nuclear charge

• The distance of the outer electron from the nucleus

• The shielding effect by electrons in filled inner shells

The effect of the size of the positive nuclear charge on ionisation energy

The greater the nuclear charge, the greater the attractive force on the outer electron and the greater the ionisation energy

The effect of the distance of the outer electron from the nucleus on ionisation energy

The force of attraction between the nucleus and the outer electron decreases as the distance between them increases. The further an electron is from the nucleus, the lower the ionisation energy

The effect of the shielding effect by electrons in filled inner shells on ionisation energy

Electrons in the filled inner shells repel electrons in the outer shell and reduce the effect of the positive nuclear charge. The more filled inner shells or subshells there are, the smaller the attractive force on the outer electron and the lower the ionisation energy

Patterns of ionisation energy across the periodic table

• Steady increase across a period due to increasing nuclear charge with no extra shielding

• Unexpected dips caused by different electron orbitals;

  • the dip between Mg and Al is caused by the p-electron being shielded by the full s-orbital

  • the dip between P and S is due to pairing of the p-electrons making the electron a little less stable (the half shell being more stable)

Shielding effect

The repulsion between electrons in different shells. Inner shell electrons repel outer shell electrons

The patterns in period 3 for ionisation energy

• The first IE increase as you go across a period due to;

  • increasing nuclear charge (more protons in the nucleus)

  • the atoms have the same number of electron shells (same shielding)

• The first IEs are lower for period 3 than period 2 due to the increased shielding of the extra full electron shell

Patterns of first ionisation energies graph

• group 3 and 6 = anomalous (slightly lower IE than expected)

Successive ionisation energies

Are a measure of the energy needed to remove each electron in turn until all the electrons are removed from an atom

Explaining successive ionisation energies

• The first electron is easiest to remove;

  • furthest from the nucleus + shielded by two full electron shells

• The second electron shows a large jump

  • this electron is in a shell closer to the nucleus and there is less electron-electron repulsion

• The next seven electrons all come from the same period and show a gradual increase in IE;

  • the p:e ratio is increasing so holds remaining electrons are held more tightly

  • greater effective nuclear charge as the same number of protons are holding fewer and fewer electrons

  • as each electron is removed, there is less electron-electron shielding and so each shell is drawn in slightly closer to the nucleus

  • the distance of the electron from the nucleus decreases, the nuclear attraction increases

• The final two electrons require much more energy;

  • they are in the electron shell closest to the nucleus

  • there is much stronger nuclear attraction and no shielding

Equations of successive ionisation energies

Charge of ion made increases by +1

Relationship between energy and frequency

E=hf

Energy = planck’s constant x frequency

Relationship between frequency in wavelength

f=c/wavelength

Frequency=speed of light/wavelength

Emission spectra

• Coloured lines on black background

• each of the lines on the emission spectra represents a drop from one energy level down to another

• energy levels converge at higher energies

• most common transition; n=2 —> n=1

• atoms are given energy, electrons are excited and the additional energy promotes then from a lower energy level to higher one

• When the source of energy is removed and the electrons leave the excited state, they fall from the higher energy level to a lower energy level

• Energy lost is released as a photon (a quantum of light energy) with a specific frequency

Absorption spectra

• lines represent electron promotions

• black lines on coloured backgrounds

• When white light is passed through the vapour of an element, certain wavelengths will be absorbed by the atoms and removed from the light

• Black lines appear in the spectrum where light of some wavelengths has been absorbed

• The wavelengths of these lines correspond to the energy taken in by the atoms to promoting electrons from lower to higher energy levels

Atomic emission spectrum of the hydrogen atom

• Consists of seperate series of lines mainly in the ultraviolet, visible and infrared regions of the em spectrum

• Atom is excited by absorbing energy, an electron jumps to a higher energy level. Electron falls back down to a lower level, it emits energy in the form of EM radiation. This emitted energy = line in the spectrum as energy of emitted radiation = difference between the two energy levels

• E=hf therefore electronic transitions between different energy levels = emission of radiation of diff freq + therefore producing different lines in the spectrum

Energy, frequency and wavelength of spectrum

Left —> right;

• energy increases

• frequency decreases

• wavelength increases

  • as frequency increases, the lines get closer together as energy difference between the shells decreases

Lyman series

• Ultraviolet region

• Transitions back to first shell/n=1 energy level (ground level)

• As frequency increases, the number of the lines on the spectrum also increases

• Highest frequency

• Highest energy

• Most common transition = n=1–>n=2

Balmer series

• visible region

• Drops back to the n=2 region

• Frequency increases

• Wavelength increases

• 1 less line than Lyman series

Paschen series

• Infrared region

• Drops back to the n=3 level

• As frequency decreases, the number of lines also decreases

Convergence limit

When the spectral lines become so close together they have a continuous band of radiation and seperate lines cannot be distinguished

Representing IE on an energy diagram

Arrow from n=1–>n=infinity

Relationship between the frequency of the convergence limit of the Lyman series and the IE of the hydrogen atom

• convergence limit represents the ionisation of the hydrogen atom

• Measuring the convergent frequency (n=1–>n=infinity) allows the IE to be calculated using E=hf