Chemistry CHEM 1002 Lecture 2 Dr Gavin Sewell gavin.sewell@tudublin.ie 1. Basic chemistry concepts, atomic Structure, electron configurations, the periodic table. 2. Bonding and intermolecular forces 3. Solid state structure 4. Gas Laws 5. Molecular shape Semester 1 Test (End of Oct/Beginning Nov): 10% January Exam: 20% Online Homework: 10% Lab: 30% Semester 2 Test: 10% May Exam: 20% For your immediate attention! There are weekly tutorials associated with this module. Attendance at tutorials is compulsory. In addition, students are required to complete a weekly activity in webcourses. It is expected that students will devote study time to practising questions and revising material between each of the lectures. Questions are available in the tutorial workbook. A Leaving Certificate chemistry book is not suitable. The recommended text book for Year 1 and Year 2 chemistry is Chemistry3 by Andrew Burrows. However, some introductory material may be more thoroughly covered in one of the following (Available in section 540 in the library). Chemistry: The Central Science, T. L. Brown, H. E. Le May and B. E. Bursten Chemistry: Molecules, Matter and Change, L. Jones and P. Atkins Chemistry, Z. V. Zumdahl States of Matter Solid – Has a rigid shape and a fixed volume that changes very little with temperature and pressure. Liquid – Has a fixed volume, but a liquid is fluid and takes the shape of its container and has no definite form of its own. Gas – The volume of a gas is not fixed, rather it is determined by the size of the container. The volume of gas varies with temperature and pressure. http://chuma.cas.usf.edu/ KMT helps us interpret the properties of solids, liquids and gases. According to this theory, all matter consists of extremely tiny particles (atoms and molecules), which are in constant motion. In solids these particles are packed closely together, usually in a regular array. In liquids and gases, the atoms or molecules are arranged at random. Liquid and gases are fluid because their particles are not confined to specific locations and can move past each other. 1. http://web2.uwindsor.ca/ 2.http://catalog.flatworldknowledge.com/bookhub/ A pure substance has a definite and constant composition-sugar, salt, water etc. It can be an element or a compound but the composition does not vary. Another feature of a pure substance is that it cannot be separated into two different species by any physical technique. If this were true our sample would be classified as a mixture. 8 An element is composed of a single kind of atom. An atom is the smallest particle of an element that still has all the properties of the element. A compound is composed of two or more elements in a specific ratio. For example, water is a compound made up of two elements, hydrogen (H) and oxygen (O). These elements are combined in a very specific way — in a ratio of two hydrogen atoms to one oxygen atom. Compounds can be decomposed to their constituent elements. Pure water can be decomposed to hydrogen and oxygen by passing an electric current through it (electrolysis). Likewise molten sodium chloride can be decomposed to yield its constituent elementssodium (Na) and chlorine gas (Cl2) 1. http://chemed.chem.purdue.edu/ 2.http://spmchemistry.onlinetuition.com Even though only 118 elements/ atoms are known, there appears to be no limit to the number of compounds that can be made from these elements. More than 20 million compounds are now known with about half a million added to the list each year. A Ruf et al., Proc. Natl. Acad. Sci. USA, 2017, https://www.chemistryworld.com/new-typeNew type of metal–organic compound of-metal-organic-compound-discovered-indiscovered in meteorites meteorites/2500483.article An atom can be defined as the smallest particle of an element that can exist while still retaining the properties of that element. A simple view of the atom is shown below. 1 http://www.aboutthemcat.org/chemistry/atomic-structure.php 2 http://web.ornl.gov/info/reporter/no4/z_con.htm Atoms are not the most elementary particles, they are comprised of subatomic particles. Thomson and Rutherford were responsible for early experiments on the nature of the atom. Joseph Thomson Ernest Rutherford Democritus 400 BCE John Dalton 1803 CE Planck, Schrödinger,de Broglie et al. 1930’s Neils Bohr 1913 Chadwick 1932 J.J. Thomson 1904 Ernest Rutherford 1911 Rutherford fired alpha particles (helium nuclei) at an extremely thin gold film. He expected them to go straight through with perhaps a minor deflection which would have been predicted by the Thomson model. Most particles went straight through, but some particles bounced directly off the gold sheet. Rutherford hypothesized that the positive alpha particles had hit a concentrated mass concentrated in the centre, rather than the diffuse positive charge proposed by Thomson. 1 http://sun.menloschool.org/~dspence/chemistry/atomic/ruth_expt.html 2 https://en.wikipedia.org/wiki/File:Geiger-Marsden_apparatus_photo.jpg Rutherford did not give complete picture: The combined mass of protons and electrons smaller than mass of atom Remaining mass is made up of neutrons – neutral particles (Chadwick, 1932) Mass of electron = 9.1 x 10-28 g (≈ 1/1837 that of a proton) Mass of proton ≈ Mass of neutron = 1.67 x 10-24 g James Chadwick 1891-1974 An element is a substance composed of only one type of atom. Atoms are composed of three subatomic particles (for our purposes at least). Particle Charge Location Mass Proton Positive (+) Nucleus 1 a.m.u Neutron Neutral Nucleus 1 a.m.u Electron Negative (-) Outside nucleus 1/1837 of a.m.u • Each element has a name and a unique chemical symbol. An element has an atomic number and a mass number. • Atomic Number: (Z) Mass Number: (A) • • The mass number is always greater than or equal to the atomic Number • The mass number gives an indication of how many protons and neutrons an atom has. • The atomic number tells us how many protons an atom has Atomic number Atomic mass How many protons, neutrons and electrons in the following: When we look at the periodic table, we see that the mass number is not a whole number. This is because elements can exist as different isotopes Isotopes are atoms that have the same atomic number but different mass number. The mass number is different as there are a different number of neutrons in the nucleus. The number on the periodic table is called the atomic mass unit (a.m.u). This depends on isotopic abundance. Radiocarbon dating relies on the radioactive decay of 14C Natural uranium is 99.3 % 238U and 0.7 % 235U Uranium is enriched for power generation and other purposes..... Calculate the atomic mass of carbon based on the following information: Isotope Abundance 12C 98.90 % 13C 1.10 % To calculate the average atomic weight, each atomic weight is multiplied by its percent abundance expressed as a decimal. The results are then added together. Carbon for instance: (12.0) (0.9890) + (13.0) (0.0110) = 12.011 a.m.u Silver (Ag) exists as two isotopes 107Ag and 109Ag. The abundance of 107Ag is 51.8% and that for 109Ag is 48.2%. (a) How many protons, electrons and neutrons are in each isotope of silver? (You will need to consult periodic table) 47 47 60 47 47 62 (b) Calculate the atomic mass unit of silver and check your answer with the periodic table. (107 x 0.518) + (109 x 0.482) = 108 a.m.u. Mass spectrometry The most accurately available method for comparing the masses of atoms involves the use of a mass spectrometer. Vapourised sample High speed Electron Beam Ionaccelerating electric field Accelerated ion beam Least massive ions Most massive ions Mass spectrometry Initially the sample is vapourised and passed into a beam of high energy electrons. These knock electrons off the atoms or molecules under examination making them into positively charged ions. These then enter an electric field which accelerates them. An accelerating ion produces its own magnetic field which interacts with an applied magnetic field. The most massive ions are deflected the smallest amount, while the least massive ions are deflected to a greater extent. The position where they hit the detector can then yield information about the relative mass of each ion. In this way different isotopes can be detected since they have differing masses. Mass spectrometry Mass spectrometry Use the data from the mass spectra shown to calculate the a.m.u of chlorine (Cl) and magnesium (Mg). Check your answers with the periodic table.
