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Chapter 2 5 September 2013 2.1 Composition of the Atom The Structure of the Atom and the Periodic Table Selected Properties of the Three Basic Subatomic Particles Name Charge Mass(amu) Mass (grams) Electrons (e) -1 5.4 x 10-4 9.1095 x 10-28 Protons (p) +1 1.00 1.6725 X 10-24 0 1.00 1.6750 x 10-24 Neutrons (n) 1 2.1 Composition of the Atom 2.1 Composition of the Atom Determining the Composition of an Atom Determine the number of protons, neutrons and electrons in each of the following: 11 5 B 55 26 Fe Isotopes • Isotopes - atoms of the same element having different masses 4 – contain same number of protons – contain different numbers of neutrons Isotopes of hydrogen Hydrogen (Hydrogen - 1) Deuterium (Hydrogen - 2) Tritium (Hydrogen - 3) 2 2.1 Composition of the Atom 2.1 Composition of the Atom Isotopic Calculations • Isotopes of the same element have identical chemical properties • Some isotopes are radioactive • Find chlorine on the periodic table • What is the atomic number of chlorine? 17 • What is the mass given 35.45 • This is not the mass number of an isotope Atomic Mass • What is this number, 35.45? • The atomic mass - the weighted average of the masses of all the isotopes that make up chlorine • Chlorine consists of chlorine-35 and chlorine-37 in about 3:1 ratio • Weighted average is an average corrected by the relative amounts of each isotope present in nature 3 2.1 Composition of the Atom 2.1 Composition of the Atom Determining Atomic Mass Calculate the atomic mass of naturally occurring chlorine if 75.77% of chlorine atoms are chlorine-35 and 24.22% of chlorine atoms are chlorine-37 Step 1: convert the percentage to a decimal fraction 0.7577 chlorine-35 0.2423 chlorine-37 Step 2: Multiply the decimal fraction by the mass of that isotope to obtain the isotope contribution to the atomic mass. For chlorine-35: 0.7577 x 35.00 amu = 26.52 amu For chlorine-37 0.2423 x 37.00 amu = 8.965 amu Step 3: sum these partial weights to get the weighted average atomic mass of chlorine: 26.52 amu + 8.965 amu = 35.48 amu 4 2.1 Composition of the Atom 2.1 Composition of the Atom Atomic Mass Determination • Nitrogen consists of two naturally occurring isotopes – 99.63% nitrogen-14 with a mass of 14.003 amu – 0.37% nitrogen-15 with a mass of 15.000 amu • What is the atomic mass of nitrogen? Ions • Ions - electrically charged particles that result from a gain or loss of one or more electrons by the parent atom • Cation - positively charged – result from the loss of electrons – 23Na 23Na+ + 1e- • Anion - negatively charged – results from the gain of electrons – 19F + 1 e- 19F- 5 2.2 Development of Atomic Theory • Dalton’s Atomic Theory - the first experimentally based theory of atomic structure of the atom. 2.2 Development of Atomic Theory Postulates of Dalton’s Atomic Theory 1. All matter consists of tiny particles called atoms 2. An atom cannot be created, divided, destroyed, or converted to any other type of atom 3. Atoms of a particular element have identical properties 6 2.2 Development of Atomic Theory 4. Atoms of different elements have different properties 5. Atoms of different elements combine in simple whole-number ratios to produce compounds (stable aggregates of atoms) 6. Chemical change involves joining, separating, or rearranging atoms 2.2 Development of Atomic Theory Postulates 1, 4, 5 and 6 are still regarded as true. Evidence for Subatomic Particles: Electrons, Protons and Neutrons • Electrons were the first subatomic particles to be discovered using the cathode ray tube 7 2.2 Development of Atomic Theory Evidence for Protons and Neutrons • Protons were the next particle to be discovered, by Goldstein – Protons have the same size charge but opposite in sign – Proton is 1837 times as heavy as electron • Neutrons – Postulated to exist in 1920’s but not demonstrated to exist until 1932 – Almost the same mass as the proton 2.2 Development of Atomic Theory Evidence for the Nucleus • Initial assumed protons and neutrons were uniformly distributed throughout the atom • Ernest Rutherford’s “Gold Foil Experiment” lead to the understanding of the nucleus – Most alpha particles pass through the foil without being deflected – Some particles were deflected, a few even directly back to the source 8 2.3 Light, Atomic Structure, and the Bohr Atom 2.2 Development of Atomic Theory Rutherford’s Gold Foil Experiment • Most of the atom is empty space • The majority of the mass is located in a small, dense region Models of the Atom (a) Thomson (b) Rutherford 9 2.3 Light, Atomic Structure, and the Bohr Atom 2.3 Light, Atomic Structure, and the Bohr Atom • Rutherford’s atom – tiny, dense, positively charged nucleus of protons surrounded by electrons • How do we describe the relationship of the electrons to each other and the nucleus? • Use the measurement of particle energy rather than position Light and Atomic Structure • Spectroscopy - absorption or emission of light by atoms. – Used to understand the electronic structure. • To understand the electronic structure, we must first understand light, electromagnetic radiation – travels in waves from a source – speed of 3.0 x 108 m/s 10 2.3 Light, Atomic Structure, and the Bohr Atom Wavelengths • Light is propagated (moves) as a collection of sine waves • Wavelength is the distance between identical points on successive waves • All wavelengths travel at the same velocity, but have their own characteristic energy Electromagnetic Spectrum high energy short wavelength low energy long wavelength 11 2.3 Light, Atomic Structure, and the Bohr Atom 2.3 Light, Atomic Structure, and the Bohr Atom Bohr Theory • Atoms can absorb and emit energy via promotion of electrons to higher energy levels and relaxation to lower levels • Energy that is emitted upon relaxation is observed as a single wavelength of light • Spectral lines are a result of electron transitions between allowed levels in the atoms • The emission emission spectrum spectrum - light of hydrogen emitted when lead to a the substance modernisunderstanding excited by an of energy the source. electronic structure of the atom 12 The Bohr Atom Electrons exist in fixed energy levels surrounding the nucleus Promotion of electron occurs as it absorbs energy Energy is released as the electron travels back to lower levels Quantization of energy Excited State Relaxation The Bohr Atom 13 Electronic Transitions • Amount of energy absorbed in jumping from one energy level to a higher energy level is a precise quantity • Energy of that jump is the energy difference between the orbits involved • Orbit - what Bohr called the fixed energy levels • Ground state - the lowest possible energy state Bohr Theory • • • • Allowed levels are quantized energy levels, orbits Electrons are found only in these energy levels Highest-energy orbits are farthest from the nucleus Atoms – absorb energy by excitation of electrons to higher energy levels – release energy by relaxation of electrons to lower energy levels • Energy differences may be calculated from the wavelength of light emitted 14 Modern Atomic Theory • Bohr’s model of the atom when applied to atoms with more than one electron failed to explain their line spectra • One major change from Bohr’s model is that electrons do not move in orbits • Atomic orbitals - regions in space with a high probability of finding an electron • Electrons move rapidly within the orbital giving a high electron density The periodic law and the periodic table • Dmitri Mendeleev and Lothar Meyer - two scientists working independently developed the precursor to our modern periodic table. • They noticed that elements have distinct regular variation of their properties when listed in order of atomic mass. • Periodic law - the physical and chemical properties of the elements are periodic functions of their atomic numbers. 15