2.1 Atoms and reactions

Transcription

1 2.1 Atoms and reactions Atomic structure and isotopes (a) isotopes as atoms of the same element with different numbers of neutrons and different masses An isotope, is defined as an atom of the same element with a different number of neutrons (and hence a different atomic mass) An isotope will, yield no change in chemical reactivity however, it may come with slight changes in the chemical properties of the element in context such as, density. For example, Silicon and it's three isotopes; 28 Si: o Protons: 14 o Electrons: 14 o Neutrons: 14 [(28) - (14)] 29 Si: o Protons: 14 o Electrons: 14 o Neutrons: 15 [(29) - (14)] 30 Si: o Protons: 14 o Electrons: 14 o Neutrons: 16 [(30) - (14)] (b) atomic structure in terms of protons, neutrons and electron for atoms and ions, given the atomic number, mass number and any ionic charge An atom, is defined as the basic unit of matter and the defining structure of the elements. Its constituents are; Protons o Relative charge on particle of (1) o Within the atomic nuclei o Made up of three quarks (two "up", one "down". Held together by a cloud of Gluons) o Relative mass of (1) Neutrons o Relative charge on particle of (0) o Within the atomic nuclei o Made up of three quarks (one "up", two "down". Held together by a cloud of Gluons) o Relative mass of (1)

2 Electrons o Relative charge on particle of (-1) o Surround the nuclei in pathways called Orbitals occupying shells (sub-shells) - also known as energy levels. o Sub-atomic particle - the structure is unknown. o Relative mass of (1/2000) Basic structure of an atom The shells of an atom fill in succession and in the pattern: 2; 8; 8; 18; 18; 32; 32 History of the Atomic structure The theory of the atom has changed vastly over the past century, swinging towards a much higher degree of accuracy, which, of course is to be expected. Scientific knowledge has developed massively over time. The development of the atomic theory: 1. Democritus developed the idea of atom (460 BC); a. His thought was that grinding solids would eventually produce pieces which were so small they could be ground no smaller. He called the "product", Atoma. 2. John Dalton (1808); a. suggested that all matter was made up of tiny spheres that were able to bounce around with perfect elasticity and called them, Atoms. 3. Joseph John Thompson (1898); a. found that atoms could sometimes eject a far smaller negative particle which he called an, Electron.

3 4. Joseph John Thompson's, 'Plum pudding model' (1904); a. Thompson develops the idea that an atom was made up of electrons scattered unevenly within an elastic sphere surrounded by a soup of positive charge to balance the electron's charge 5. Ernest Rutherford (1910); a. oversaw Geiger and Marsden carrying out his famous experiment. They fired Helium nuclei at a piece of gold foil which was only a few atoms thick. They found that although most of them passed through. About 1 in 10,000 hit. Some helium nuclei deflected whereas, some bounced directly back at them. b. These recent findings, aided Rutherford to propose a new model. One with a central nucleus. He suggested that the positive charge was all within a central nucleus. Thus, holding the electrons in place by means of electrical attraction. 6. Niels Bohr (1914) a. Bohr refined Rutherford's idea by adding that the electrons were in orbits. Rather like planets orbiting the sun. With each orbit only able to contain a set number of electrons.

4 Basic atom composition/sub-atomic particle calculations Basics: 1. Number of Neutrons = (Mass Number) - (Atomic number). 2. Number of Protons = Atomic number. 3. Number of Electrons = Number of protons or atomic mass of an element. For example, Carbon; 1. Neutrons: 12 [(12) - (6)] 2. Protons: 6 3. Electrons: 6 Ions Atoms, don't always have the same number of protons and electrons. When this happens, it will set the relative charge of the atom above or below, (0). This is what's called an ion - essentially being a charged atom. Often, atoms need to Borrow or Donate electrons to another atom, to allow bonding to occur. This is known as, ionic bonding. The atom/s donating the electron, gain a positive charge. The atom/s receiving the electron, yield a negative charge. For example, Sodium Chloride (NaCl);

5 In the example above, the chloride atom, borrows 1 electron from the Sodium atom and this makes the overall charge of the atom (-1). This is due to the factuality that, the number of protons (+ve subatomic particle) the number of electrons (-ve subatomic particle) hence, the relative charge does not equal 0. Likewise, with the Sodium atom: there are now less electrons than protons, as the sodium atom donates 1 electron to the outermost shell of the chloride atom - the atom has a charge of (+1). Knowing this, you should be comprehend why the composition of an atom changes, if the atom has an overall relative charge not equal to zero. Here are some examples of ions; Mg 2+ : o Protons: 12 o Electrons: 10 [(12) - (2)] o Neutrons: 12 H - : o Protons: 1 o Electrons: 2 [(1) + (1)] o Neutrons: 1 N 3- : o Protons: 7 o Electrons: 10 [(7) + (3)] o Neutrons :7 Note: The sign i.e. - or +, must be wrote after, the integer for the relative charge of the atom - in context of ions. (c) explanation of the terms relative isotopic mass (mass compared with 1/12th mass of carbon-12) and relative atomic mass (weighted mean mass compared with 1/12th mass of carbon-12), based on the mass of a 12 C atom, the standard for atomic masses Scientists, need to have a method to compare the masses of elements, to help them quantify the number of atoms, in a sample, for example. Clearly, they won't be using the common units of mass I.e. Kilograms or grams, all measurements are calculated in atomic mass units (a.m.u). This is due to the factuality that the mass of protons & neutrons are so small, and not to forget they're almost the same, so the resolution of common units simply isn't great enough to differentiate between the two sub-atomic particles - essentially, it increases precision. The use of the unit, enables chemists to construct a relative scale. 1 a.m.u, is equal to, 1/12 of the mass of a Carbon-12 atom. Hence, a Carbon-12 atom (isotope of Carbon) is equal to 12 a.m.u. However, the relative atomic mass of elemental carbon is slightly larger at, a.m.u. The reasoning behind this being that, multiple isotopes of the element naturally occur in a sample of Carbon i.e. Carbon-12, Carbon-13 and Carbon-14. Terminology (& simplified definitions); The relative atomic mass (A r) of an element, is the weighted mean mass of its atoms, taking into account the relative proportions of each isotope present in a naturally occurring sample of an element, on a scale that assigns carbon-12 a relative mass of exactly 12. o The weighted mean mass of an atom of an element compared to the mass of 1/12 of the mass of an atom in carbon 12.

6 The relative isotopic mass, is the mass of an atom of the isotope, on a scale where a carbon-12 atom has a relative mass of exactly 12. o The mass of an atom of an isotope compared to the mass of 1/12 of the mass of an atom in carbon-12 (d) use of mass spectrometry in: (i) the determination of relative isotopic masses and relative abundances of the isotope Mass spectrometer, stages; 1. Vaporize the sample 2. Ionize the sample. (Think: how are ions formed? By the 'loss' of an [ ]?) 3. Focussing and accelerating the ions 4. Deflecting the ions 5. Detecting the ions Ionizing the sample is essential, as it allows for the isotopes to separate further after vaporization and hence, direct the ion to the 'detector' to be analyzed, in a more singular approach. Typically, the ions have a charge of (+1). The above image is a mass spectra (the graph that mass spectrometers plot the data on) of a sample of Boron. The two peaks represent the two relative isotopes present in the sample. The x-axis showing the title m/z, essentially represents the relative isotopic mass of the isotopes in the sample, those in this particular case being; Boron-10 (mass of 10) and Boron-11 (mass of 11) -- the number after the elemental name is the relative mass of the isotope. The y-axis represents, the relative abundance, and hence, the amount of atoms in Boron of that specific isotope. In this case, Boron-10 has a relative abundance of 23 and Boron-11, of 100. The x-axis (m/z) = mass-to-charge ratio. The charge is dependant on how many electrons are knocked off in the initial process (I.e. mass spectroscopy) - in this case it was 1, so the ionic charge equals (+1). m (mass) / z (charge) = relative mass of ion.

7 (ii) calculation of the relative atomic mass of an element from the relative abundances of its isotopes Using the example/data from above: Considering the information provided above, hypothetically speaking, if we had 123 atoms of Boron, 23 of those atoms would be Boron-10 and 100 of the atom would be Boron-11. To calculate the relative atomic mass of Boron is relatively straightforward, once you've picked out the essential information from the graph. It is as follows; (relative atomic mass) = (Relative abundance x isotopic mass) / (sum of abundances of isotopes) 1. Relative atomic mass = [Boron-10 (23 x 10)] + [Boron-11 (100 x 11)] / ( ) 2. Relative atomic mass = ( ) / (123) 3. Relative atomic mass = 10.8 (3 s.f) (e) use of the terms relative molecular mass, M r, and relative formula mass and their calculation from relative atomic masses. Relative molecular mass & relative formula mass, both are calculated exactly the same, whereas their definitions vary in terms of the specifics, however. Relative molecular mass (M r): the sum of all the atomic masses for all the atoms in a given formula of a simple molecule. o Example, N 2; M r = M r = 28 Relative formula mass: the sum of all the atomic masses for all the atoms in a given formula of a compound that has a giant molecular structure. o Example, NaCl; M r = M r = 58.5