Atomic Structure and the Periodic Table

Atomic Structure and the Periodic TableYou should be able to:1.1 discuss the process of theoretical change with respect to Daltons atomic change;1.2 describe the structure of the atom;

1.3 define the following terms: (i) mass number; (ii) isotopes; (iii) relative atomic and isotopic masses based on the carbon-12 scale. Specific Objectives1.4 explain the phenomenon of radioactivity;

1.5 cite the use of radioisotopes;1.6 calculate the relative atomic mass of an element, given isotopic masses and abundances;1.7 explain how data from emission spectra provide evidence for discrete energy levels within the atom;Specific Objectives1.8 describe the atomic orbitals;

1.9 describe the shapes of the s and p orbitals;

1.10 determine the electronic configurations of atoms and ions in terms of s, p and d orbitals;1.11 state the factors which influence the first ionisation energy of elements;

Specific Objectives1.12 explain how ionisation energy data provide evidence for sub-shells;

1.13 derive the electronic configuration of an element from data on successive ionisation energies.Examples of criteria that are considered when theories are acceptedFit between evidence and theoretical constructs

Reliability and accuracy of data

Replicability of experiments

Examples of criteria that are considered when theories are acceptedConsensus within the scientific community

Societal factorsElectronic Structure Summary of Subatomic particles

Note that the accurate masses of the protons and neutron are not exactly the same. The charges on the proton and electron are exactly the same. NameMass numberMass/kgRelativechargeCharge/coulombProton11.673 x 10-27+1+1.6 x 10-19

Neutron11.675 x 10-27

00electron00.911 x 10-30

-1-1.6 x 10-19

RADIOACTIVITYA large number of elements have unstable atoms which split to form smaller particles. In the process energy is released. This energy is in the form of radiation.

RADIOACTIVITYThree types of radiation are given off by radioactive substances. They differ in their response to an electric field. The uncharged rays, (gamma) rays, are similar to X rays. They have high penetrating power, being able to pass through 0.1 m of metal. RADIOACTIVITYMeasurement of m/z identified (alpha) rays as the nuclei of helium and (beta) rays as electrons. rays can pass through o.o1 m of metal, and rays can penetrate no more than 0.o1mm of metal. RADIOACTIVITY- Nuclear equationsAlpha-decayWhen an atom disintegrates by losing an alpha-particle (), the remaining fragment will have a mass number four less than the original atom and an atomic number two units lower.

RADIOACTIVITY- Nuclear equationsRemember that both mass (indicated by the superscripts) and charge (indicated by the subscripts) must be balance in nuclear equations. Most radioactive isotopes with atomic numbers greater than 83 (bismuth) undergo apha-decay. CLASSWORKWrite nuclear equations for the alpha-decay of radium-226 () and plutonium-238 (), use the periodic table to find the symbols of the breakdown products. RADIOACTIVITY- Nuclear equationsBeta-decayElements with an atomic number less than 83 do not usually exhibit alpha-decay. Instead, various isotopes of these elements emit beta-particles. For example, carbon-14 () decays by beta particle emission forming nitrogen (). The decay process can be represented by the equation: (See Chemistry in Context- page 78)RADIOACTIVITY- Nuclear equationsBeta-decayDuring beta-decay, mass number remains constant, so the number of protons + neutrons stays constant. However, the atomic number increases by one. CLASSWORKWrite nuclear equations for the beta-decay of strontium-90 () and Iodine-131 ().

How do chemical reactions differ from nuclear reactions such as alpha-decay and beta-decay? Rate of radioactive decayThe rate at which a radioactive isotope decays cannot be speeded up or slowed down. It depends only on the identity of the isotope and the amount of isotope present. The rate of radioactive decay is proportional to the number of radioactive atoms present. Such reactions are described as first-order reactions. The time taken for radioactive isotopes to decay to half the number of radioactive atoms is called the half-life, t1/2.

Stability of IsotopesThe time taken for an isotope to break down into a new atom gives an indication as to how stable the isotope is. The more stable the isotope the longer or greater its half life. The rate of natural decay of a radioactive isotope is not affected by physical or chemical means. The half life of radioactive isotopes vary from fraction of a second to millions of years, e.g. Cobalt have a have life of 5.2yrs. NOTE TO STUDENTSRead up on gamma ()-rays. Read up and make notes on the use of radioisotopes. When atoms react, electrons are redistributed. These electrons may be transferred from one atom to another or they may be shared between the reacting atoms in a different wayIt is possible to obtain information about the arrangement of electrons in atoms by study the ease with which atoms lose electrons. The energy needed to remove one electron from an atom is known as the ionisation energy (ionisation enthalpy) and is given the symbol Hi QuestionWhat is the first ionisation energy of an element? ANSWERThe energy required to remove one electron from each atom in a mole of gaseous atoms producing one mole of ions with one positive charge.

Thus, the first ionisation energy of sodium is the energy required for the processNa(g) ------> Na+(g) + e- H = Hi1The most useful method of determining the ionisation energies of elements involves a study of their emission spectraThe compounds of some elements emit light when heated. For example, the flame colour of lithium compounds is red and that of potassium compounds is lilac. Some gaseous elements will also emit light when they are subjected to large potential differences in electric discharge tubes. Neon advertising signs work in this way and the yellow sodium street lamps are discharge tubes containing sodium vapour. What is a spectroscope?A spectroscope is a device which splits light into its component wavelengths or frequencies. Line Spectra If the light emitted by these substances is examined using a spectroscope, separate lines of different colour will be observed. This kind of spectrum is called a line emission spectrum. Each element has its own characteristic set of lines different from any other element. Line Spectra Elements can be identified from their line emission spectra. Line SpectraEach line in an emission spectrum corresponds to light of a particular frequency. Formation of Line SpectraElectrons in an atom can exist only at certain energy levels. Under normal conditions, the electrons in an atom or ion fill the lowest energy levels first. Formation of Line SpectraWhen sufficient energy is supplied to the atom, it is sometimes possible to promote (excite) an electron from a lower energy level (the ground state) to a higher one. This process is called excitation. Formation of Line SpectraThe electron is unstable in the higher energy level, so it will emit the excess energy as radiation and drop back into a lower energy level. Formation of Line SpectraThe energy difference between the higher and lower energy levels can have only certain fixed values because the energy themselves are fixed. Formation of Line SpectraAt each drop in level, it gives out a quantum of energy equal to the difference in levels, E.Because of the relationship E = h, each quantum has a different frequency and therefore contributes one line to the line spectrum. Formation of Line SpectraThe radiation emitted when an electron falls from a higher to a lower energy level can have only certain fixed frequencies (i.e. Certain specific colours) because the frequency of any radiation is determined by its energy. Emission SpectraThe small amount of radiation emitted by an electron when it falls from a higher to a lower energy level is referred to as a quantum of radiation. Electrons and OrbitsElectrons occupy only certain fixed energy levels around a nucleus. These energy levels can be visualised as orbits of increasing radius. Orbits of larger radius have higher energy. In the hydrogen atom, which has only one electron, only one energy level is occupied at any time. The situation is rather like a set of shelves with one single object. The object can only be placed on a shelf (not in the gaps between) and the shelves still exist even if they have no object on them. THE BOHR MODELIn 1913, Niels Bohr put forward his picture of the atom. Bohr referred to Max Plancks recently developed quantum theory, according to which energy can be absorbed or emitted in certain amounts, like separate packets of energy, called quanta. Bohr suggested:1. An electron moving in an orbit can have only certain amounts of energy, not an infinite number of values: its energy is quantised. THE BOHR MODEL (Contd)2. The energy that an electron needs in order to move in a particular orbit depends on the radius of the orbit. An electron in an orbit distant from the nucleus requires higher energy than an electron in an orbit near the nucleus.3. If the energy of the electron is quantised, the radius of the orbit also must be quantised . There is a restricted number of orbits with certain radii, not an infinite number of orbits. THE BOHR MODEL (Contd)4. An electron moving in one of these orbits does not emit energy. In order to move to an orbit farther away from the nucleus, the electron must absorb energy to do work against the attraction of the nucleus. If an atom absorbs a photon (a quantum of light energy), it can promote an electron from an inner orbit to an outer orbit. If sufficient photons are absorbed, a black line appears in the absorption spectrum. Energy and FrequencyThe relationship between the energy (E) of a quantum of radiation and its frequency () is:E = h X h is a constant, called Plancks constant. The value of h is 4 x 10-13 kJ s mol-1

BRAIN EXERCISEWhat is the value of Plancks constant with unit kJ s molecule-1 ? HYDROGEN SPECTRUM The electronic energy levels are numbered (n=1, n=2, n=3, etc.). The numbers are sometimes referred to as the principal quantum numbers for the energy levels which corresponds to the shells of electrons. THE HYDROGEN SPECTRUMThe colours in the visible region of the hydrogen spectrum are caused by electron transitions from higher levels to the n=2 and not the lowest energy level (n=1).

Lines which result from transitions to the n=2 level in an atom are known as the Balmer seriesLines which result from transitions to the lowest energy level in an atom are known as the Lyman Series THE HYDROGEN SPECTRUMAs the energy levels get closer and eventually come together, it follows that the spectral lines also get closer and eventually come together. If sufficient energy is given to an atom, it is possible to excite an electron just beyond the highest energy level. In this case the electron will escape and the atom becomes an ion. Ionisation has taken place. In an atom, the highest possible energy level corresponds to the frequency at which the lines in the spectrum come together. So, by determining the frequency at which the converging spectral lines come together, the ionisation energy of an element can be determined. This particular frequency is called the convergence limit. CLASS ACTIVITYFrequencies of the Lyman Series for hydrogen

Frequency / 1o 14 HzTransition to which frequency corresponds24.66n=2 to n=129.23n=3 to n=1 30.83n=4 to n=131.57n=4 to n=1

31.97n=4 to n=1

32.21n=4 to n=1

32.37n=4 to n=1

CLASS ACTIVITYWork out the difference in frequency () between successive lines in the Lyman Series for hydrogen.Plot a graph of (vertically) against frequency, . (Use the value of the lower frequency for plotting .)Use your graph to estimate the frequency when becomes 0.

4. becomes 0 when difference in energy between the electronic energy levels becomes 0. Use the relationship E=h to find the energy which corresponds to the frequency when becomes 0. This energy is the ionisation energy (ionisation enthalpy) for hydrogen.An accurate value for the frequency at the convergence limit for hydrogen is 32.7 x 1014 Hz. Using E=h, the energy of radiation with this frequency= 4 x 10-13 x 32.7 x 1014 kJmol-1 = 1308 kJmol-1

Using ionisation energies to predict electronic structures- evidence for shellsIf an atom containing several electrons is provided with sufficient energy it will lose one electron. Additional supplies of energy will result in the removal of a second electron, then a third, then a fourth, and so on. A succession of ionisations is therefore possible, each of which has its associated ionisation energy. For example, the first ionisation energy of sodium corresponds toNa(g) ------> Na+(g) + e- H = +494kJmol-1

Whereas the second ionisation energy of sodium corresponds toNa(g) ------> Na+(g) + e- H = +4564kJmol-1

(see page 67 of chemistry in context)

Electrons occupy atomic orbitals in pairs.In each pair, the electrons are spinning in opposite directions. Chemist believe that paired electrons can only be stable when they spin in opposite directions so that the magnetic attraction which results from their opposite spins can counterbalance the electrical repulsion which results from their identical negative charges. Evidence for sub-shells of electronsBy studying ionisation energies and atomic spectra, scientists have concluded that: the n=1 shell can have 2 electrons in the same sub-level (sub-shell) the n=2 shell can have 2 electrons in one sub-level, and 6 electrons in a slightly higher sub-level. the n=3 shell can have 2 electrons in one sub-level, 6 electrons in a slightly higher sub-level and 10 electrons in a still slightly higher sub-level Evidence for sub-shells of electrons (Continuation)The n =4 shell can have 2 electrons in one sub-level, 6 electrons in a still slightly higher sub-level, 10 electrons in a still slightly higher sub-level, and 14 electrons in a still slightly higher sub-level. Evidence for sub-shells of electrons (continuation)The sub-shells (or sub-levels) containing 2 electrons are called s sub-shells (s electrons). The sub-shells (or sub-levels) containing 6 electrons are called p sub-shells (p electrons).The sub-shells (or sub-levels) containing 10 electrons are called d sub-shells (d electrons). The sub-shells (or sub-levels) containing 14 electrons are called f sub-shells (f electrons)When energy sub-levels are being filled, electrons always occupy the lowest available energy sub-level first and the electrons pair-up as soon as each sub-level is filled.The electron structure of an atom can be described in terms of its sub-shells occupied by electrons . A number (1, 2, 3, 4 etc.) is used to denote the quantum shell, a letter (s, p, d or f) to denote the sub-shell and a superscript to indicate the number of electrons in the sub- shell. Electron ConfigurationThe electronic structure for argon can be written in terms of energy levels as 2,8,8; and more precisely in terms of energy sub-shells as 1s2 , 2s2 , 2p6 , 3s2 , 3p6

Note: The 4s sub-shell is lower in energy than the 3d sub-shell and thus is occupied first. Electron Configuration Calcium, with 20 electrons, would be:1s2 , 2s2 , 2p6 , 3s2 , 3p6 , 4s2Sometimes it simplifies things to refer back to the previous noble gas structure. So the electron arrangement of calcium, Ca, could be written [Ar] 4s2 as a shorthand for 1s2 , 2s2 , 2p6 , 3s2 , 3p6 , 4s2 , since 1s2 , 2s2 , 2p6 , 3s2 , 3p6 is the electron arrangement of argon.

The filling of the 3d sub-shells is very important in chemistry of the elements from Sc to Zn. These are known as the transition elements (commonly referred to as the d-block elements. Electrons and orbitalsOrbitals are regions in which there is the greatest probability of finding particular electrons, although the electrons are not confined to these regions. Each orbital can hold either one or a maximum of two electrons. If the orbital contains two paired-up electrons they will be spinning in opposite directions. Electrons and OrbitalsThe following simple rules are followed when writing electronic configuration:1. Aufbau principle: Orbitals are filled so that those of lowest energy are filled first. (Aufbau is German for building up.)2. Pauli exclusion principle: no two electrons in an atom can have the same four quantum numbers. In other words, only two electrons may occupy the same atomic orbital, and these must have opposite spins. Electrons and Orbitals3. Hunds rule: When we come to orbitals of equal energy (degenerate orbitals) such as the p orbitals, we add one electron with their spins unpaired until each of the degenerate orbitals contains one electron. (This allows the electrons, which repel each other, to be farther apart.) Then we begin adding a second electron to each degenerate orbital so that the spins are paired.