Survey
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
2. Atomic structure “A wrong theory is always so much better than no theory at all.” William Lawrence Bragg (1890-1971) ● Richard Feynman (1918-1988): “If all of scientific knowledge were to be destroyed, and only one sentence passed to the next generation....I believe it is that all things are made of atoms.” Daltons atomic theory 1. All matter consists of very small particles called atoms. 2. An element consists of atoms of one type only. 3. Compounds consist of atoms of more than one element and are formed by combining atoms in whole-number ratios (the law of multiple proportions). 4. In a chemical reaction atoms are not created or destroyed. Thomsons “plum-pudding” model of the atom Rutheford´s gold foil experiment ● This experiment tested Rutherfords “plum-pudding model” and led to the conclusion that the atom is mainly empty space. http://www.youtube.com/watch?v=wzALbzTdnc8&feature=results_main&playnex t=1&list=PLDF89CB07DD4E004E The atom In a neutral atom, the number of electrons equals the number of protons. Sub-atomic particle Charge Mass (kg) Mass/amu proton +1 1,67 ▪ 10-27 kg 1 neutron 0 1,67 ▪ 10-27 kg 1 electron -1 9,1 ▪ 10-31 kg 0,0005 ● ● Mass number (A) = number of protons plus neutrons in the nucleus of an atom Atomic number (Z): the number of protons in the nucleus of an atom. Isotopes ● ● ● Atoms of the same element always contain the same number of protons (= same atomic number Z), but if they contain a different number of neutrons (= different mass number A) they are called isotopes. Most elements occur in nature as a mixture of isotopes. For most of the light elements, the numbers of protons and neutrons in the nucleus are nearly equal. Radioisotopes ● Are used in: - Nuclear medicine for diagnostics, treatment and research (iodine-131, iodine-125, cobalt-60) - Tracers in biochemical and pharmaceutical research - Chemical clocks in geological and archaeological dating (carbon-14) https://www.youtube.com/watch?v=udkQwW6aLik Half-life, t1/2 Radioactive decay • Many isotopes of elements are radioactive, because their nuclei break down spontaneously and emit radiation: • alpha, α: - two protons and two neutrons - can not penetrate the skin • beta, β: - Stream of electrons that comes from the nucleus, NOT from the electron shells • gamma, γ: - High energy electromagnetic radiation - Extremely penetrating - Blocked efficiently by very dense materials: Pb Relative atomic mass • The mass of an atom depends on the number of protons and neutrons in the nucleus (the mass of electrons is so small that it can be ignored in chemistry). ● ● Since the mass of a single atom is very small (on average 10-23 g), it is not possible to weight single atoms or molecules. The relative atomic mass of an element is the weighted mean mass of all the naturally occuring isotopes of that element relative to the mass of carbon-12. The unified atomic mass unit ● ● One-twelfth of the mass of a carbon-12 atom in its ground state. 1 amu or 1 u = 1.6605402 x 10-27 kg The relative atomic mass, Ar ● The ratio of the average mass of the atom to the unified atomic mass unit. The mass spectrometer • The relative atomic mass of an element can be determined using a mass spectrometer: 1) the sample is vaporized 2) ionized to positive ions: M (g) → M+ (g) + e3) the ions are accelerated in an electric field 4) the ions are deflected by an external magnetic field 5) the ions are recorded on a detector (relative amounts of the ions, mass to charge ratio) https://www.youtube.com/watch?v=mBT73Pesiog 2.2 Electron configuration Metals often give different characteristic colours when heated strongly, i.e. given energy. e.g. Cu: Ba: Na: Li: ● Each element has its own characteristic colour which can be used to identify the element. ● The electromagnetic spectrum (EMS) • Visible light is one type of electromagnetic radiation. ● The peak to peak distance is called the radiation's wavelength, λ. • The number of waves which pass a point in one second is the frequency of the radiation, v. • Wavelength is related to the frequency of the radiation: c = vλ c = speed of light ≈ 3,00 · 108 m s-1 v = frequency of the radiation (Hz) λ = wavelength (m) Continuous spectrum ● When white light (sunlight) is passed through a prism, a continuous spectrum of all colours can be obtained. ● A continuous spectrum contains light of all wavelengths and so appears as a continuous series of colours. Line spectra ● ● When white light is passed through hydrogen gas, some of the light is absorbed. An absorption spectrum is produced where some colours are missing ( those that are absorbed by hydrogen). ● A corresponding emission line spectrum can be produced. The lines correspond to the light of particular wavelengths given off (=emitted) by the element. http://www.youtube.com/watch?v=nM5Kg7RUoTE The hydrogen spectrum ● How can a hydrogen atom absorb and emit energy? • Niels Bohr proposed a theoretical explanation: ● ● electrons travel in orbits around the nucleus The energy of each orbit is fixed (=quantized) http://www.youtube.com/watch?v=nM5Kg7RUoTE ● ● Light can be described as a stream of photons or tiny “packets” (= quanta) of light energy. The energy of electromagnetic radiation is expressed: E = hv E = energy of a photon (J) h = Planck's constant ≈ 6,626 · 10-34 J s v = frequency of the radiation (Hz) ● The smaller the wavelength (= the higher the frequency), the greater the energy of the photon. Excitation • When an atom absorbs energy, an electron moves into an orbit of higher energy further from the nucleus. An unstable excited state is produced. • The electron soon falls back to a lower level and gives out (= emits) the energy in the form of electromagnetic radiation with a specific frequency, v. • The lowest energy level of the electron is called the ground state. ∆Eelectron = Ef- Ei = hv ● The line spectrum of hydrogen provides evidence for the existence of electrons in discrete energy levels, which get closer together (=converge) at higher energies. Electron arrangement ● ● In the Bohr model of the atom the energy levels are drawn as shells. Each shell or orbit can be occupied by a certain number of electrons. Electron configuration • The Bohr´s model of an atom is a simplification that only explains the spectral lines of hydrogen. • Limitations: • http://www.youtube.com/watch?v=ofp-OHIq6Wo The quantum mechanical model of the atom • de Broglie: "electrons behave like particles and have the properties of waves“ • Heisenberg: "it is impossible to calculate the exact location of an electron at an exact moment in time = the uncertainty principle“ • Schrödinger: "Schrödingers wave equation describes the region around a nucleus where there is a high probability of finding electrons“ http://en.wikipedia.org/wiki/Schr%C3%B6dinger_equation ● ● According to the QUANTUM MECHANICAL model the electron has to be treated as both a particle and a wave. It is impossible to know exactly where an electron is at any given time, but it is possibly to calculate the region in which it is likely to be found. • Atomic orbital: A region in space around an atomic nucleus in which there is a high probability of finding the electron. Energy levels • Each electron in atom can be described by four different quantum numbers: • The principal quantum number describes the main energy level, n= 1, 2, 3, etc. Maximum number of electrons 2n2. • The second quantum number describes the type of orbital in that level: l= 0,1,2,3 or s,p,d,f. • The shapes of s, p and d orbitals /watch?v=sMt5Dcex0kg ● ● The third quantum number determines the number of orbitals of each type in each level. For example p x, py, pz, which are localized along the x, y and z-axes respectively. The fourth quantum number describes the spin of the electron around its own axis, either clockwise or counter-clockwise, s= +1/2 and -1/2 or ↑↓. Sub-levels of electrons Writing electron configurations • The electron configuration of an element describes how the electrons are distributed in sub-levels. Orbital diagram ● In order to show the spin of each electron, boxes can be used to represent orbitals. Aufbau Principle: • electrons are placed into orbitals of lowest energy first. Pauli exclusion principle: • no two electrons in an atom can have exactly the same four quantum numbers no more than two electrons can occupy any one orbital, and if two electrons are in the same orbital they must have the opposite spin = each orbital can contain a maximum of two electrons 1s Hund´s rule: ● If more than one orbital in a sub-level is available, for example px, py, pz , electrons occupy the orbitals single with parallel spin before they are paired up = orbitals within the same sub-shell are filled singly first Ex: Apply the electron-in-a-box to determine the electron configuration of aluminium. Al (Z = 13) Full electron configuration: 1s22s22p63s23p1 Condensed electron configuration: [Ne]3s23p1 Ex: Apply the electron-in-a-box to determine the electron configuration of zink. ● ● Electron configurations can be written with the principal quantum number followed by the letter of the sub-level, with a subscript to indicate the number of electrons in that sub-level. e.g. Na: 1s22s22p63s1 When writing electronic configurations, the sum of the superscripts must always total the number of electrons in the atom (or ion). IB requires that you can write the configuration for any element or ion up to krypton (Z=36). Halv-filled and filled sub-levels are more stable: Cr Cu Halv-filled and filled sub-levels are more stable: Cr Cu Electron configuration of ions ● Positive ions are formed by the loss of electrons. ● Negative ions are formed by the gain of electrons. ● (An electron in a doubly occupied orbital is repelled by its partner and therefore easier to remove than an electron from a half-filled orbital.) ● Worked example: Apply the electron-in-box method to determine the electron configuration of the calcium-ion and the chlorine-ion : Transition metals ● ● ● Elements with partially filled d sub-levels are called transition metals. When the d-block elements form positive ions the energy levels are attracted more strongly to the nucleus. The 3d sub-level drops below the 4s sub-level in energy and therefore the 4s electrons are removed first. ● Worked example: Apply the in-box method to determine the electron configuration of the Fe -ion and the copper-ion : Evidence from ionization energies Ionization energy (kJ mol-1) • The energy required to remove one electron from an atom in its gaseous state is called the first ionization energy of the element. X (g) → X+ (g) + e• The second and third ionization energies would therefore correspond to the chemical equations: X+ (g) → X2+ (g) + eX2+ (g) → X3+ (g) + e- Patterns in first ionization energies Evidence for sub-levels ● ● Successive ionization energies for the same element can also be measured. Graph of successive ionization energies for sodium. ● ● ● ● As more electrons are removed, the pull of the positively charged nucleus holds the remaining electrons more tightly. Increasingly more energy is required to remove them. The graph does not increase regularly, which provides evidence that the main levels are split into sub-levels. Identify the sub-levels from which the electrons are removed. Ionization energies from emission spectra ● ● The hydrogen emission spectrum consists of a series of lines that converge at higher energies. These lines are associated with electron transitions from upper energy levels back down to lower levels. ● ● When an electron falls from the limit of convergence (n=∞) and returns to the ground state (n=1) energy is given out. This energy corresponds to the first ionization energy. Ex. When the hydrogen electron falls from n=∞ to n=1 it emits energy in the form of light with the wavelength 91,2 nm. This produces a line in the UV region of the hydrogen emission spectrum (Lyman series). Calculate the first ionization energy (kJ mol -1) for hydrogen. E = hv c = vλ E = energy of a photon (J) h = Planck's constant = 6,63 · 10-34 J s v = frequency c = speed of light ≈ 3,00 · 108 m s-1 v = frequency λ = wavelength (m)