Chapter 2 Atoms, Molecules, and Ions. Part 2

Group Activity • Draw you idea of what an atom “looks like”. Chapter 2 • Answer the following questions: – How do you know what an atom “looks like” at the atomic level? – How do scientists know what at atom “looks like”? – Can scientists see the parts of an atom? Explain. Atoms, Molecules, and Ions Conservation of Mass Before chemistry there was… Alchemy 1774: Priestley isolated oxygen by heating HgO (mercury (II) oxide). When carried out in a closed container the total mass of the contents within the container didn’t change. 1785: Lavoisier formulated The Law of Conservation of Mass. In a chemical reaction, matter is neither created nor destroyed Mass of reactants must = mass of products + any unreacted material 6g H2 8g H2 Alchemy’s demise 1661: Robert Boyle, the father of modern chemistry, publishes The Sceptical Chymist, Law of Definite Proportions Joseph Louis Proust discovered that all samples of a specific substance will have the same ratio of elements by mass. For example, all samples of pure: •water: 1 part H, 8 part O by mass •carbon dioxide: 1 part C, 2.7 parts O by mass Elements within a compound combine in specific proportions Compounds are not a random mixture of elements + 34g O2 42g total reaction 36g H2O 42g total Dalton’s Atomic Theory • Dalton's Theory:
All matter is composed of atoms; indivisible particles that are exceedingly (really, really) small. All atoms of a given element are identical, both in mass and in chemical properties. However, they are different from atoms of other elements. Atoms are not created or destroyed in chemical reactions. Reactions only change how atoms are arranged. Atoms combine in small, whole-number ratios to form compounds. 7 Plum Pudding? Nuclear Atom Based on J.J. Thompson’s experiments, the “plum pudding” model of the atom was developed • Gold foil experiment: Positive alpha particles were “shot” to toward gold foil. - - - - - • In 1911, Hand Geiger and Ernest Marsden carried out an experiment under the direction of Ernest Rutherford. (Radium) - Negatively charged electrons scattered throughout Large sphere of positive charge 10 What was expected vs. what was observed Expected data: no large deflections. Plum pudding model fits. Actual data: large deflections. New model of the atom needed. 1 out of every 8000 α particle had a large deflection - - - - - ? - + - - Discovery of the Neutron • 1932! • Sir James Chadwick (Irene Joliot-Curie, Frederic Joliot-Curie) • Bombarded beryllium with α particles • The radiation that was emitted was electrically neutral • This radiation could eject protons from wax • Used multiple elements • Determined it had the same mass as a proton Discovery of the neutron Can we see atoms? Can we see inside them? • Simplified experimental set-up • Why did it take so long to discover the neutron? STM image (7 nm x 7 nm) of a single zigzag chain of cesium atoms (red) on a gallium-arsenside surface (blue) 4 4 Number Nickel atoms – STM Relative Sizes • Nucleus is to an atom what a pea is to a stadium Atomic #, Mass #, and Isotopes Atomic Mass vs. Atomic Mass Number • Atomic mass and atomic number are NOT the same How many of each particle are in the following atoms (protons, electrons, neutrons)? 1 H+ 1 2 H 1 3 H 1 12 6C 13 6C 14 6C 18 8O 17 8O 16 8O 14 7N 15 7N 238 U 92 3- • Atomic mass is the mass found on the periodic table, which is an average of isotopes. This mass is not a whole number. It has units. • Isotopic mass is the mass of a specific isotope. Must look up from a source. Except for one important exception (12C), it is not a whole number. It has units. 2- • Mass number is the number of protons plus the number of neutrons for a specific isotope. This MUST be a whole number! It has no units. 15 O 8 234 U 92 O 16O 16O Atomic mass Isotopic mass Mass number 15.999 amu 15.994 amu 16 21 Covalent Chemical Bonds Molecules, Ions, and Chemical Bonds • Which side of the periodic table gains electrons? Loses electrons? – Metals tend to lose electrons, have a positive charge (cation) – Nonmetals tend to gain electrons, have a negative charge (anion) 1+ 2+ • • • • Na  Na1+ Cl  Cl1Mg  Mg2+ S  S2- 3- 2- 1- • Covalent bonds form when atoms share electrons; typically between nonmetals • Hydrogen, oxygen and other elements exist as diatomic molecules – H2, N2, F2, O2, I2, Cl2, Br2 Representations of ionic compounds Ionic Chemical Bonds • Ionic bonds form when atoms transfer electrons; typically between metal and nonmetal • Do they appear different than covalent compounds? • Ions are atoms or groups of atoms that have a positive or negative charge; they have gained or lost electrons. Anions have gained one or more electrons; since electrons are negative, anions have negative charges Cations have lost one or more electrons; cations have positive charges Ionic vs. Covalent Bonding Metallic Bonding • Can you determine which is most likely ionic bonding and which is covalent bonding? Terminology Compounds H2O Atom: one sphere in an image Electrons are shared by all of the metal atoms. The are said to be delocalized electrons. Sometimes referred to as a “sea of electrons”. Molecule: individual unit made up of two or more atoms covalently bonded • Compound: composed of one or more atoms of different elements bonded covalently or ionically. Are the following compounds? Molecules? Formula Unit: individual unit made up of two or more ions NaCl bonded O2 O3 NaCl Classifications of Matter • Pure substance: all atoms, compounds, or molecules are the same in a sample • Describe each picture below using the terms – mixture or pure substance – atom, element, molecule, and/or compound • Mixture: blends of two or more different substances – Homogeneous: uniform throughout (seawater) – Heterogeneous: not uniform (water and oil) Physical Separation Chemical Separation 32 Classifications of Matter Chemical Formulas • How would you classify the following? • Chemical (or molecular) formulas indicate the composition of molecules and ions – Which atoms are present and the ratios in which they have combined – Air – Distilled water – Ice Cream – Rocks – Gold jewelry – Diamond • Structural formulas show how atoms are bonded together • Empirical formulas give the smallest whole-number ratio of a molecular formula. (H2O2 becomes HO.) These formulas are often used in analysis. 33 Chemical Formulas • Why do we have different types of models? Sodium Chloride NaCl ---------- 35