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Atomic Structure Chapter 2 Knowing the Atom Composed of subatomic particles Protons (p+) Neutrons (n0) Electrons (e-) Consists of 2 regions Nucleus Electron shells History of Chemistry The atomism concept is thousands of years old and was originally founded for philosophical reasoning The earliest references to atomism date back to sixth century BC India Later emerged with Leucippus and Democritus in Ancient Greece History Cont. John Dalton developed the atomic theory in 1803 explaining why chemical elements react in simple proportions by mass He proposed that each element had its own set of atoms that bond to form compounds Did not prove atoms existed, but he did show his atomic theory was consistent History Cont. Avogadro made the distinction between atoms and molecules (O vs. O2) in 1811 Was not clarified and accepted until 1858 due to the work of Cannizzaro Dalton’s atomic theory was modified throughout the twentieth century and his original symbols for the elements replaced by our modern symbols developed by Berzelius, a Swedish chemist Relative mass and charges of subatomic particles Protons have positive charge Neutrons have no charge These two have almost the same mass (relative mass= 1) Electrons have a negative charge and very small relative mass (5E-4) The opposite electrical charges hold the atom together Relative mass and relative charges of subatomic particles Neutron help stabilize the nucleus by separating the protons and reducing electrostatic repulsion If too many or too few neutrons are present the nucleus will undergo radioactive decay Electrons control the chemical properties of the elements Each element has a unique set of chemical properties History of Chemistry Electron discovered in 1897 by J.J. Thompson, a British physicist He did a series of experiments with cathode ray tubes and cathode rays to discover them He found electron mass to be 1000 times less than that of a hydrogen ion He also demonstrated that hydrogen has only a single electron History of Chemistry Thompson also developed a model of the atom known as the “Plum Pudding Model” This model was later changed to the atomic model we rely on now by Ernest Rutherford Mass Number, Atomic Number and Isotopes Mass Number The total number of protons and neutrons (Symbol A) Atomic Number The number of protons in the nucleus (Symbol Z) Isotope Atoms of the same element that have different mass numbers due to a different amount of neutrons AZX Notation Notation practice Write the following elements in AZX notation Sodium Hydrogen Aluminum Oxygen Ions An element forms an ion when one or more electrons are added or removed from the atom A positive ion is formed with electrons are removed (cation) A negative ion is formed when electrons are added (anion) To determine the number of electrons in an atom, add or subtract from the number of protons Isotopes Many elements exits as a mixture of isotope Each has a unique mass and abundance Used to calculate average atomic mass Isotopes have identical chemical properties, but slightly different physical properties Ex. Uranium-235 and uranium-238 (forms 99.3% natural uranium) History of Chemistry French chemist Henri Becquerel discovered uranium salts (released radiation which passed through the wrapping paper around a photographic plate) Further experimented by Pierre and Marie Curie, who named it radioactivity The Curies also discovered the elements radium and polonium Radioactivity Many elements contain unstable nuclides which break up spontaneously with the emissions of ionizing radiation The unstable nuclides are described as being radioactive and are called radioisotopes Radioisotopes Radiation Relative Charge Relative Mass Nature Penetration Deflection by electric field Alpha particles +2 4 2 protons and 2 neutrons (He2+ ion) Stopped by a few sheets of paper Low Beta particles -1 1/1837 Electron Stopped by a few mm of plastic or Al High Gamma rays 0 0 Electromag Stopped by netic a few cm of radiation of lead very high frequency None Radioactive Decay The rate which nuclei undergo radioactive decay varies between elements Radioactive decay is an exponential process Rates are compared using half-life (the time taken for half of the radioactive nuclei to undergo decay) During alpha and beta decay a more stable isotope forms Radioactive Decay Each radioisotope has a unique half-life Unaffected by temperature or pressure Ex. Iodine-131 has a half life of 8 days Due to half-life, different radioisotopes are used to diagnose and treat diseases Ex. Cobalt-60, a powerful gamma emitter, used to treat different types of cancer (Radiotherapy) The Mass Spectrometer Allows chemists to accurately determine the relative atomic masses of atoms Can also be used to determine relative molecular masses and structures of compounds Look at Interactive syllabus Determining Relative Atomic Mass Relative atomic mass=(% isotope A * mass isotope A)+ (% isotope B *mass isotope B)… Ex. Chlorine- 35 and chlorine -37 have percent abundances of 75% and 25% Rubidium exists as a mixture of two isotopes 85Rb and 87Rb. The percentage abundances are 72.1% and 27.9%. The relative atomic mass of gallium is 69.7. gallium is composed of two isotopes: gallium-69 and gallium-71. Calculate the percentage abundance of gallium-69. Light as Waves and Particles Light is emitted in the form of electromagnetic waves These consist of an oscillating electric wave and magnetic wave which travel in a sinusoidal pattern The two waves are arrange perpendicularly Light Wave Terminology Wavelength (λ) The distance between two neighboring crests or troughs of a wave Frequency (ν) The number of waves which pass a point in one second. Units are hertz (Hz); 1/sec. Speed (c) The distance travelled by a wave in one second. Units are m s-1. Wave equation c=λν where c= 3.0E8 m s-1 Terminology Photons “packets” of light energy. The term used when light is described by a particle model Planck’s constant (h) A value of 6.63E-34 J s. Used in the equation: E=hν. Continuous spectrum The spectrum of colours formed from white light which is composed of visible light of a certain range of wavelengths History of Chemistry Issac Newton was the first western scientist to show that sunlight is composed of many colours. He used a prism to split the light into a spectrum then re-formed it into white light (Late 1600s- early 1700s) A Muslim scientist, Ibn al-Haytham, performed similar light dispersion experiments during the Golden Age (965-1040) Electron Excitation When gaseous atoms are excited, they emit light of certain wavelengths Excitation is when an electron is raised to a higher level, then light is emitted when it returns to the unexcited state Could be a thermal or electrical process Thermal a flame is formed Ex. Fireworks exploding Electrical a high voltage is passed across a gaseous sample of an element at low pressure Ex. Neon signs Emission Spectra When light from atoms with excited electrons are passed through a prism, emission spectra are formed These spectra are called line spectra Each chemical element has its own unique spectrum which could be used to identify that element Line Spectra Line spectra differ from continuous spectra in two ways: 1. An emission (line) spectrum is made up of separate lines (discontinuous) 2. the lines converge, becoming progressively closer as the frequency or energy of the emission lines increases History of Chemistry Johann Jakob Balmer (1825-1898), Swiss mathematics teacher who studied visible lines of the hydrogen emissions spectrum Developed a mathematical formula to calculate wavelengths His formula allowed for discovery of additional lines on the hydrogen spectrum and of white stars Neils Bohr proposed a model in 1913 to explain why Balmer’s formula gave correct wavelengths Energy Levels and Spectra Bohr suggested a theory to account for the complete emission spectrum of helium He thought electrons moved in orbits where they had fixed amounts of potential energy The further an electron from the nucleus, the greater the energy Bohr’s theory An electron moving in an orbit does not emit energy Energy Levels In order to move into an orbit further from the nucleus: The electron must ABSORB electric or thermal energy Do work against the attraction of the positive nucleus The electron is then in the excited state Electron Excitation and Spectra Electrons dropping back from higher energy orbits for emission spectra The energy of the light is equal to the difference between the two energy levels: ΔE=hf Energy Levels Bohr labeled his energy levels with n and a number n=1 is the lowest energy level (nearest the nucleus) Considered “Ground state” The most stable state for a hydrogen atom The next highest energy state is n=2, etc. If an electron receives enough energy, it is completely removed from the atom; forming an ion The amount of energy required to remove the electron is the ionization energy (transition from n=1 to n=∞) Hydrogen Atom H Absorption and Emission Hydrogen Photon Energy Calculate the energy of photons that five rise to the red line of the Balmer series. The wavelength is 656.3 nm and the speed of light in a vacuum is 3.0 x 108 m s-1. f= c/λ E=hf (h=6.63 x10-34 J s) History of Chemistry Neils Bohr (1885- 1962) a Danish physicist awarded Nobel Prize in physics (1922) for atomic structure contributions Also part of Manhattan Project; developed first atomic bomb Studied with Rutherford and J.J. Thompson He determined that electrons can only have certain values or quantities of energy His theory could only account for the spectra of hydrogen, thus his theory was flawed and later replaced by Shrödinger’s quantum mechanical model Electron Arrangement Electrons are arranged in electron shells Hydrogen has only a single electron and it is located in the shell nearest the nucleus The first shell (first energy level) can hold a maximum of 2 electrons The second shell (second energy level) can hold a maximum of 8 electrons The third shell (third energy level) can hold a maximum of 18 electrons Electron Arrangement Indicates the number of electrons without drawing the shells Ex. Lithium has 3 electrons and an arrangement of 2,1 or 2.1 Ex. Sodium has an arrangement of 2,8,1 or 2.8.1 Ionization Energy (HL) This force is dependent on two main factors and is fine tuned by a third factor. Size of the atom The nuclear charge Shielding effectElectrons between the outer electron and the nucleus (fine tuning) Atomic Radius Distance of the outer electrons from nucleus As distance between electron and nucleus increases, the attraction decreases Causes ionisation energy to decrease Nuclear Charge As nuclear charge becomes more positive, the attraction to electrons becomes stronger Increases ionisation energy Shielding Effect Outer valence electrons repelled by inner electrons and kept from the pull of the nucleus Mostly effective if electrons are close to nucleus Electrons in the first shell have the highest shielding effect The more shielded from the nucleus, the lower the ionization energy HL Ionizing Energies Comparing the I.E. of H and He If we compare the energy required to remove the outermost electron of hydrogen (element number 1) and helium element number 2) we see that approximately twice the energy is needed for helium. As the nuclear charge of helium is twice that of hydrogen we can conclude that the electron is roughly the same distance from the nucleus in both cases. The first ionisation energy The first ionisation energy is defined as the energy required to remove one mole of electrons from one mole of gaseous atoms to provide one mole of gaseous single charged ions. Na(g) --> Na+(g) + 1e Subsequent ionisation energies are defined in a similar way only by removing electrons from already charged ions. The second ionisation energy Na+(g) --> Na2+(g) + 1e Successive electrons can be stripped from an atom until there is only the nucleus left. If the energy required to achieve this for each ionisation is plotted on a graph (with a log scale) against the ionisation number, the 'jumps' in the required energy clearly show the main and sub energy levels. Ionization energy trend Drops in I.E. There are dips in ionization across periods 2 and 3 First decrease is because electron shielding is changed due to loss of sub-level electron (Between Be and B) More E required for Be to lose a lower-level electron Second decrease between N and O Easier to remove electron from O because it is part of a pair; N has a single electron in that orbital Hund’s Rule Every orbital in a sub-shell is single occupies with one electron before any one orbital is doubly occupied Reason for the drop between N and O Practice using Hund’s rule with: Nitrogen N+ Oxygen O+ O2- Relating I.E. and Configuration I.E. are always endothermic: energy absorbed and work done to remove electron(s) Second I.E. always larger than the first Further successive ionization energies increase Successive Ionisation E of K Interpreting the graph Electron arrangement is 2.8.8.1 Which part of the graph shows the first electron(s) removed? After which removed electrons are the largest increases in I.E.? What do these large jumps correspond to? The y-axis is measured in log10 to reduce the values of the “jumps” Construct a Graph Sketch a graph of the successive ionization energies of silicon against electrons removed Start by determining the electron arrangement for Si Then, start from the outer electrons and work your way to the inner Working in Reverse The first 8 experimentally determined I.E. for a chemical element are as follows (kJ mol-1) 580 1800 2750 11580 14850 18400 23300 27500 Deduce the following: Number of e- in outer shell of an atom of the element The group in the periodic table which the element belongs Outer detailed configuration of the atom Orbitals and Energy Levels Electrons are found in shells around nucleus, numbered 1, 2, 3 etc Each shell has sub-shells, labeled s, p, d, or f Number of sub-shells are equal to the number of the shell Shell (energy level) number Number of sub-shells in the shell Sub-shells (sub-levels) 1 1 1s 2 2 2s and 2p 3 3 3s, 3p, 3d 4 4 4s, 4p, 4d, 4f Sub- Shells Each sub-shell contains a particular number of orbitals, where the electrons are located Number of orbitals depends on the type of sub-shell Each orbital can hold a maximum of 2 electrons Type of sub-shell (sub-level) Number of orbitals Maximum number of electrons in the subshell (sub-level) s 1 2 p 3 6 d 5 10 f 7 14 Energy shells split into sub-shells Diagram • Note that the 4s sub-shell has lower energy than 3d, so the electrons will fill it before the 3d Electron Configurations Write a complete electron configuration for the first 15 elements on the periodic table (H P) Shapes of Orbitals The four types of orbitals all have different shapes s orbitals are spheres and differ in size and energy p have a dumb-bell shape and have different orientations (px, py, pz) and do not differ in energy (degenerate) History of Chemistry Erwin Schrödinger developed Schrödinger wave equation to give energy levels of H atom and the shapes and energies of orbitals; published in 1926 Werner Heisenberg proposed that it was impossible to measure both location and momentum of an electron; Heisenberg’s uncertainty principle published in the mid-1900’s A scientist can measure the position with some accuracy, but the momentum has a large range and vice versa Filling Atomic Orbitals Pauli exclusion principle each orbital can hold a maximum of two electrons Aufbau principle electrons occupy an empty atomic orbital with the lowest energy Hund’s rule within a sub-shell electrons experience repulsion, thus enter two different orbitals of the same energy Electrons can spin clockwise or anticlockwise Single electrons in sub shells have the same spin Electron Configurations of Atoms Hydrogen: 1s 1s1 Lithium: 1s 1s22s1 Helium: 1s Carbon: 1s22s22p2 1s2 1s 2s 2s 2p Electron Configurations to Know! (HL) 1-38 49-54 Electron Configurations for Ions Electron configurations of ions are determined based on the same principles as neutral atoms The added or lost electrons are simply included or excluded from the arrangement Ex: Mg 1s22s22p63s2 Mg2+ 1s22s22p6 Octet Rule Atoms gain or lose electrons to achieve a complete outer electron orbital with 8 electrons The noble gases have a complete orbital, thus are unreactive