Particle Symbol Charge / C Relative charge Mass / g Relative mass. Electron e x x x 10-4 (negligible)

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1 ATOMIC STRUCTURE Atom An atom is the smallest part of an element which can ever exist. Atoms consists of a small central region called nucleus in which all the mass of the atom (protons and neutrons) are concentrated. This nucleus is surrounded by a much larger volume in which the electrons move. The nucleus of an atom is about one ten-thousandth of the size of the whole atom. If we magnified an atom to the size of a football stadium, the nucleus would be represented by a pea placed at the centre of the pitch. Particle Symbol Charge / C Relative charge Mass / g Relative mass Proton p x x Neutron n x Electron e x x x 0-4 (negligible) Any nucleus is characterized by two numbers : Z A X where Z is the atomic number = no. of protons A is the mass number = no. of protons + neutrons Atomic number is more significant because : all elements have their own unique atomic numbers 2 all atoms of the same element must have the same number of protons but their mass may be different from one another as a result of the existence of isotopes (atoms of the same element with different number of neutrons) Therefore, elements in the Periodic Table are arranged in the order of atomic numbers. Relative atomic mass is the ratio of the weighted mean isotopic mass of the atom to 2 of the mass of a 2 C atom. At first, the element hydrogen was chosen as the standard against which the masses of other atoms were compared. It was because chemists realized that it had the smallest atoms which could conveniently be assigned a relative atomic mass of one. Later, when relative atomic masses could be obtained more accurately, carbon-2 was chosen as the new standard because : carbon is a very common element; 2 being a solid, it is easier to store and transport than hydrogen, which is a gas. Calculate the relative atomic masses for argon and potassium using the following data : Ar : 36 (0.34) ; 38 (0.06) ; 40 (99.6) K : 39 (93.2) ; 40 (0.2) ; 4 (6.76) MASS SPECTROMETER Relative masses of atoms and of molecules can be accurately determined by a mass spectrometer. An analogy to the mass spectrometer would be to roll a bowling ball and a basketball at the same speed at a target while a stiff crosswind is blowing. The basketball is lighter, therefore, its path is more readily changed by the crosswind.

2 Atom 2 component function A to vaporize the sample (which may be an element or a compound) the sample is bombarded by fast moving electrons to form positive ions B C D E X(g) + e - X + (g) + 2 e - fast slow to accelerate the beam of positive ions to deflect the ion beam so that ions of a particular mass/charge (m/e) ratio are focused into the ion-detector (the lighter the positive ions, the greater the deflection) to detect the signal and pass it on to a recorder By varying the strength of accelerating electric field or deflecting magnetic field, ions of any m/e ratio can be brought to the ion detector. A mass spectrum showing the m/e ratio of the ions and the corresponding intensity (i.e. the relative amount of that ion) can then be traced out by a recorder. In most mass spectra, the values of the m/e ratio can be converted to the relative masses of the particles if the charges on the ions are taken to be one. Each peak in the mass spectrum represents the relative abundance of a particular type of particle with a certain isotopic mass. (No instrumental details and mathematical treatments are required) The following mass spectrometer trace was obtained for a naturally occurring sample of an element X Relative abundance Relative isotopic mass (i) Give an interpretation of this trace. (ii) Calculate the relative atomic mass of the naturally occurring sample to 4 sig. fig.

3 Atom 3 The mass spectrum of chlorine molecule contained three peaks at m/e 70, 72 and 74 of relative intensity 9 : 6 : respectively. Deduce the relative atomic mass of chlorine. The material to be analyzed may be an element or a compound : when sample is an element, the peaks in the mass spectrum can give information about various isotopes of the element 2 when sample is a compound, the peak with the highest m/e ratio will most likely correspond to the molecular ion, i.e. the molecule which has lost only a single electron Sketch the mass spectrum for (i) methane, CH 4 ; and (ii) hydrogen chloride, HCl. (Assuming H, 2 C and 4 N have 00% abundance among their isotopes) ELECTRONIC STRUCTURE OF ATOMS Our knowledge of the way in which electrons are distributed in an atom comes from two sources of evidence : study of emission spectrum of atomic hydrogen 2 study of ionization energies of the elements Emission Hydrogen Spectrum When an electrical discharge is passed through hydrogen at low pressure, a coloured glow is observed. Upon dispersion by a prism, the radiation can be spread out into a few sharp, coloured lines, which are commonly known as the emission hydrogen spectrum.

4 Characteristics of Emission Hydrogen Spectrum Atom 4 It consists of several series of discrete lines in different parts of the electromagnetic spectrum; e.g. Lyman series in the ultraviolet (UV) region, Balmer series in the visible region and Paschen series in the infra-red (IR) region 2 It consists of coloured lines against a dark background 3 The lines in each series converge (get closer) towards higher frequencies (or shorter wavelengths), until they eventually merge into a continuum 4 Every element has its own unique pattern of lines in its emission spectrum frequency, ν / 0 4 Hz Balmer worked out a relationship to describe the pattern of lines in the visible spectrum : ν α n 2 The lines in all series are found to fit in the general formula : = R ( λ n - m ) 2 2 where λ is the wavelength of the particular line R is Rydberg s constant ( =.097 x 0 7 m - ) n, m are integers which, in turn, have the following set of values : Lyman series n = m = 2, 3, 4, 5... in the UV region Balmer series n = 2 m = 3, 4, 5, 6... in the visible region Paschen series n = 3 m = 4, 5, 6, 7... in the IR region The wavelength of the first line in the Lyman series is found to be.25 x 0-7 m. (a) Calculate the wavelength of the third line in the Balmer series. (b) Sketch the first five lines in the Balmer series in the above emission spectrum. m st line 2 nd line 3 rd line 4 th line 5 th line λ λ / 0-7 m ν / 0 4 Hz

5 Interpretation of Emission Hydrogen Spectrum Atom 5 To account for the hydrogen spectrum, it is assumed that electrons in an atom is quantized (i.e. electrons can only move in orbits with a certain fixed amount of energy). Under normal conditions, the electrons in an atom or ion fill the lowest energy levels first (the ground state). Upon absorption of sufficient energy (e.g. via electrical discharge), it is possible to promote (excite) an electron from a lower energy level to a higher one (the excited state). The electron is now unstable in the higher energy level, so it will emit the excess energy as radiation and drop back into the lower energy level. The energy difference between the higher and lower energy levels can have only certain fixed values because the energy levels themselves are fixed. This means that transition of an electron from a higher to a lower level emits a photon (discrete amount of energy in the form of radiation) with energy equal to the difference in energy between the two energy states, hence produces a discrete line in the emission spectrum. The spectrum lines in each series converge towards higher frequencies as a result of the fact that the energy difference between orbits decrease with increase in principal quantum number (i.e. higher energy levels). The relationship between the energy (E) of radiation and its frequency (ν) is E = h ν where h is Planck s constant ( = 6.63 x 0-34 J s) The atoms of each element have a unique arrangement of electrons with definite energy levels. When the atom is excited, electrons are brought to higher energy levels. As the electrons return to lower energy levels, the atom emits radiation of definite wavelengths. This results in a unique emission spectrum which can provide useful information about the atom and serve as an identification for the element concerned. Ionization Energies If sufficient energy is applied to an atom to excite an electron from its ground state to just beyond the highest possible one, then the electron can escape from the atom. The atom is said to be ionized. 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, we can find the ionization energy of an element. This particular frequency is called the convergence limit. The wavenumber of the first line in the Balmer series is found to be.525 x 0 6 m -. Calculate the ionization energy per mole of hydrogen atoms.

6 Successive Ionization Energy - Evidence of Shells Atom 6 If 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, a fourth, and so on, i.e. a succession of ionizations is possible. The successive ionization energies of beryllium are 900, 760, 4900 and 2060 kj mol - respectively. (i) Write equations to represent the succession of ionization energies of beryllium. (ii) I.E. Sketch the trend for the succession of ionization energies of beryllium Each successive ionization requires a progressively larger amount of energy than the preceding one because : every time an electron is being removed, the resulting particle carries one more positive charge and exerts a stronger electrostatic attraction on that electron 2 electrons are being removed from a progressively lower energy levels, which are much closer to the nucleus and thus being held more tightly A plot of succession of ionization of an element gives information about its electronic arrangement (e.g. 2, 2 in the case of Be). It also proves that two of the four electrons spend most of the time closer to the nucleus than the other two. They are said to occupy different electron shells. Sketch the successive ionization energies of sodium and carbon. Evidence of Sub-shells st I.E atomic number

7 Trends in the plot of st I.E. of the first twenty elements : Down the group, I.E. decrease one more shell distance between nucleus & outermost electron increase more readily to remove(ionize) outermost electron 2 Across the period, I.E. increase size decrease outermost electron being held by nucleus more tightly more difficult to remove outermost electron 3 Within each period, I.E. is minimum at Group I and maximum at Group 0 atoms tend to have their outermost shell full-filled in order to attain an extra stability 4 Within each period, broken trend is found in a pattern of 2, 3, 3 suggest the presence of sub-shells based on similar reason as in (3) Atom 7 Electronic Configurations Atomic Orbitals Louis de Broglie postulated that matter could have both particle and wave properties. This postulate was later confirmed by the experiment that demonstrated the diffracting property of electrons, proving that electrons did indeed possess wave properties. Electrons, therefore, are not localized in any fixed orbits. We cannot know exactly where an electron is at any particular moment, but we can only describe the probability of finding the electron in a certain position at any time. The region in space where an electron is likely to be found is called an orbital, and is best illustrated by a charge cloud diagram where the dot density represents the probability of finding the electron. Suppose a photograph of the electron of hydrogen could be taken at any instant, and a second photograph taken at an instant later when that electron would occupy a different position. Millions of such photographs superimposed would generate an electron density diagram, the best representation of an atomic orbital. Atomic orbitals have no definite boundary owing to the probability description. Instead, their shapes are more significant properties in describing atomic orbitals : s-orbital -- spherically symmetric (i.e. variation is exactly the same in any direction from the nucleus) -- electron cloud density is densest near the centre, but falls away as the distance increases from the nucleus -- electron density is zero at the nucleus and at infinite distance from the nucleus p-orbital -- dumbbell-shaped -- each p-orbital (namely p x, p y and p z ) has its axis mutually perpendicular to each other d-orbital -- 4 out of 5 orbitals have 4 lobes extending out perpendicular to each other -- the last one has 2 lobes extending along z-axis with a doughnut-shaped ring around the centre of x-y plane The average distance of s electrons from the nucleus is less than that of p electron, so that s electrons are more firmly held by the nucleus. d xy d xz d yz p x p y p z d x 2 y 2 d z 2

8 Building Up Electronic Configurations Rules in building up electronic configurations : Aufbau s building up principle electrons will enter the orbitals in order of increasing energy 2 Hund s rule orbitals of the same energy must be occupied singly before pairing occurs 3 Pauli s exclusion principle no orbitals can accommodate more than TWO electrons Notation in writing electronic configurations : main shell s no. of electrons 2 in the orbital type of sub-shells / orbitals quantum shells n = 4 n = 3 n = 2 n = Atom 8 subshells 4d 4p 3d 4s 3p 3s 2p 2s s Electronic Configurations from H to Kr atomic number element symbol notations using s, 2s, 2p, etc. electron-in-boxes diagram hydrogen helium H He s s 2

9 Atom 9 atomic number element symbol notations using s, 2s, 2p, etc. electron-in-boxes diagram Some important points to note : There is a division of the Periodic Table into four different areas, namely the s-, p-, d- and f- blocks, depending on the type of the outermost electron s 2s 2p 2 4s sub-shell always has a top priority than 3d sub-shell in filling up or removing electrons 3 Chromium and copper exhibit abnormal configurations in order to attain an extra stability from the half-filled and full-filled 3d sub-shell respectively 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f Summary of Electronic Structure of Atoms A study of : emission spectrum of atomic hydrogen proves that electrons can exist in a certain discrete energy levels 2 succession of ionization for a particular element leads to its electronic configuration 3 the trend of first ionization energies introduces the presence of sub-shells within main shells