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Atomic Structure. Simple model of an atom. An atom is made of a tiny nucleus with electrons orbiting around it. The nucleus is made up of protons and neutrons. Much of an atom is empty space. The protons. Each proton has a +1charge. Each proton has a mass of 1 atomic mass unit.

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Atomic Structure


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    1. Atomic Structure

    2. Simple model of an atom • An atom is made of a tiny nucleus with electrons orbiting around it. • The nucleus is made up of protons and neutrons. • Much of an atom is empty space

    3. The protons • Each proton has a +1charge. • Each proton has a mass of 1 atomic mass unit. • The number of protons in an atom is called its atomic number(z).

    4. The neutrons • The neutrons have no charge. • Each neutron has a mass of 1 atomic mass unit. • The total number of protons and neutrons in an atom is called its mass number (A).

    5. The Electrons • Each electron has a charge of -1. • Electrons have negligible mass of 1/1840 that of a proton.

    6. Convention for writing an atom A zX

    7. Isotopes • Isotopes are atoms of the same element which have the same number of protons but different number of neutrons.

    8. Relative Isotopic Mass • Relative Isotopic Mass= Mass of one isotope of the element/(1/12) X Mass of one atom of 126C. • An atom of 12C has a relative atomic mass of 12 exactly.

    9. Relative Atomic Mass(Ar) • Ar = Weighted average of the isotopic masses/(1/12) X mass of one 126C atom

    10. Calculating the relative atomic mass of chlorine • Chlorine has two isotopes 35Cl and 37Cl in relative proportions of 75% and 25% respectively. • The weighted average mass of a chlorine atom is 35X(75/100) + 37 X (25/100) = 26.25+9.25=35.50(no unit)

    11. Relative Molecular Mass(Mr) • Mr = Mass of one molecule/(1/12)X Mass of one 126C atom. • The relative molecular mass can be worked out by adding the relative atomic masses of all the atoms present in one molecule.

    12. Mass spectrometry • A mass spectrometer separates the isotopes of an element according to their masses and shows the relative numbers of the different isotopes present. • Before the isotopes can be separated, they must be converted to positive ions.

    13. The workings • Evacuation of the instrument • Vaporisation of liquid or solid samples • Production of positive ions • Acceleration of positive ion • Deflection of positive ions • Detection of positive ions according to mass(m)/charge(e). When charge =+1, mass/charge = mass

    14. The uses of a mass spectrometer • To find the isotopic composition of an element. • To work out the relative atomic mass of an element. • To find the relative molecular mass and the fragmentation pattern of a molecule. • In forensic science.

    15. The mass spectrum of Cl % Abundance • Chlorine has two isotopes: 35Cl and 37Cl with relative proportions of 75% and 25% respectively. 75 25 35 37 m/ e

    16. First ionisation energy • It is the energy required to remove one electron from each of one mole of gaseous atoms to form one mole of gaseous ions with single positive charge and one mole of electrons. • The equation for first ionisation energy of element A is: A(g) A+(g) + e

    17. Successive ionisation energies • Successive ionisation energies provide evidence for the existence of quantum shells or electronic energy levels.

    18. Successive ionisation energies • If an atom has two electrons, it will have two ionisation energies, first ionisation energy and second ionisation energy. • If an atom has three electrons, it will have three separate ionisation energies. • All these ionisation energies for each element are its successive ionisation energies.

    19. The successive ionisation energy graph of Be • The diagram indicates two electronic energy levels. • Electrons 1 and 2 are at a higher energy level • Electrons 3 and 4 at a lower energy level -nearest to the nucleus. Log IE/kJ mol-1 x x x x 1 2 3 4 Ionisation no

    20. Electron configuration- key points • Each element has a characteristic emission spectrum which can be used to identify it. • The electrons in an element can exist only at certain energy levels - shells and sub-shells • The region in which an electron moves for most of the time is called an orbital. • An orbital can hold two electrons.

    21. The Line Spectrum • An electron can absorb sufficient energy and move to a higher energy level. • When such an electron drops to a lower energy level, the energy absorbed is given out. • The amount of energy given out appears as a line in the line spectrum of the element.

    22. First ionisation energies of successive elements - H toNe • These provide evidence of shells and subshells. • The first shell can have one sub-shell, s subshell. • The second shell can have two subshells, s and p 1st IE/kJ mol-1 x x x x x x x x x 1 2 3 4 5 6 7 8 9 10 z

    23. The aufbau principle • Electrons always occupy the lowest available energy sub-level or subshell. • Electrons pair up after a sub-level is half filled. • Numbers 1, 2, 3 denote the shells. Letters s, p, d, f denote the subshells. A superscript indicates the number electrons. • Sequence of energy levels: 1s2s2p3s3p4s3d

    24. Subshells and orbitals • An s sub-shell has only one orbital. • A p sub-shell has three orbitals. • A d subshell has five orbitals.

    25. Sub-shells and electrons • An s sub-shell can have a maximum of two electrons. • A p sub-shell can have a maximum of six electrons. • A d sub-shell can have a maximum of ten electrons.

    26. Shape of an s orbital • An s orbital is spherical in shape • The sphere is made up of a cloud of negative charge from the electrons

    27. Shape of p-orbitals - dumb-bell shaped • The three p orbitals, px,, py and pz. py P=x pz px

    28. Electron configuration of elements • H 1s1 He 1s2 • Li 1s2 2s1 Be 1s22s2 • B 1s22s22p1 C 1s22s22p2 • N 1s22s22p3 O 1s22s22p4 • F 1s22s22p5 Ne 1s22s22p6 • Na 1s22s22p63s1 Mg 1s22s22p63s2