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Chemical Foundations: Elements, Atoms, and Ions

Chemical Foundations: Elements, Atoms, and Ions. Chapters 4&5 Pages 85-155. Early History of Chemistry- 500 BC. Greeks (around 500-400 BC) divided matter into four elements: earth, air, fire, and water.

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Chemical Foundations: Elements, Atoms, and Ions

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  1. Chemical Foundations: Elements, Atoms, and Ions Chapters 4&5 Pages 85-155

  2. Early History of Chemistry- 500 BC • Greeks (around 500-400 BC) • divided matter into four elements: earth, air, fire, and water. • Aristotle- Continuous theory of matter matter can be divided and subdivided into smaller and smaller parts indefinitely- each part, no matter how small would retain the original properties of matter. • Democritus – Discontinuous theory of matterbelieved all matter is made up of tiny partices, or ATOMOS (indivisible), which could not be broken down into smaller particles

  3. History- 500 BC1600s • Alchemy- the belief that cheaper metals could be turned into more valuable metals such as gold. • Alchemists discovered many elements in their attempts to convert metals to gold. Many elements’ names have Latin origins

  4. History- 1600s Robert Boyle (1627-1691)- • Science should be grounded in experiments • Termed ELEMENTS- a substance was an element unless it could be broken down into two or more simpler substances

  5. Key Concepts and Vocabulary: Atoms, Elements and Molecules • Atom: smallest unit of an element • Element: substance that is composed of one type of atom. There are114 elements on the periodic table- 88 are naturally occurring, the remainder have been made in laboratories. • Molecule: a collection of atoms chemically bonded. • Molecules can be made from only one element, such as H2 or O2 • Molecules can be made from different elements, such as H2O or CO2

  6. The Complexities of Molecules • English language- composed of alphabet containing 26 letters. • Elements are the fundamental building blocks of matter in the universe. Distribution of the most abundant elements in Earth’s crust

  7. The Complexities of Molecules One segment of a hemoglobin molecule

  8. History- 1800’s • Most natural materials are mixtures of pure substances. • Pure substances are either elements or combinations of elements called compounds. • Law of Constant Composition-

  9. Dalton’s Atomic Theory • 1766-1844- English School Teacher • Elements are made up of tiny particles called atoms. • All atoms of a given element are identical • The atoms of a given element are different from those of any other element. • Atoms of one element can combine with atoms of other elements to form compounds. A given compound always has the same relative numbers and types of atoms. • Atoms are indivisible in chemical processes. That is, atoms are not created or destroyed in chemical reactions. A chemical reaction simply changes the way the atoms are grouped together.

  10. Impact of Dalton’s Theory

  11. Formulas and Compounds • Compound-

  12. Chemical Formula- • Elements’ symbols are used to simplify the formulas. Subscripts are used to indicate the number of each type of element. Glucose- C6H11O6

  13. History- 1800s –1900sThe Structure of the Atom • JJ Thomson (1897) Cathode Ray Tube- Credited with discovering the electron.

  14. History- 1800s –1900sThe Structure of the Atom • If negative charges are present and the overall charge on an atom is neutral, Thomson theorized that a positively charged particle must also exist. William Thomson (Lord Kelvin)- Plum Pudding Model of the Atom

  15. History- 1900sThe Structure of the Atom • Ernest Rutherford (1911)- Gold Foil Experiment • Alpha Particles- + charged particles 7500 times more massive than electrons.

  16. History- 1900sRutherford’s Gold Foil Experiment

  17. History- 1900s • The Results of Rutherford’s Gold Foil Experiment led scientists to believe that the atom must have a densely packed positive center (NUCLEUS) around which tiny electrons moved in a space that was otherwise empty. • 1919- Rutherford concludes that the atom’s nucleus must have particles positively charged called protons and that the number of protons in an atom equal the number of electrons.

  18. History- 1900sThe Nuclear Atom • Rutherford’s collegue- James Chadwick determined that the mass of the atom did not equal the mass of electrons + protons. Therefore, there must be another particle without a charge (NEUTRON) found in the nucleus.

  19. Introduction to the Modern Concept of Atomic Structure • From a very simplistic perspective- the atom is composed of a tiny nucleus (10-13 cm in diameter) and electrons that move about the nucleus at a distance of about 10-8 cm from it. • If the nucleus was a golf ball, the electrons would be about 2 Km away. • The nucleus contains protons and neutrons. • The mass of the protons is about equal to the mass of neutrons. • Neutrons have zero charge, electrons –1, protons +1. • The number of protons = the number of electrons in a neutral atom.

  20. X = the symbol of the element • A = mass number (number of protons + neutrons) • Z = atomic number (number of protons)

  21. p e The Hydrogen Atom • One electron orbiting a nucleus • Z = atomic number • N • Total mass = • Singly ionized Hydrogen is missing one electron = 1H+ 1H

  22. p p n n e e The Helium Atom • Two electrons orbiting a nucleus • Z = atomic number • N • Total mass = • Singly ionized Helium is missing one electron =4He+ • Doubly ionized Helium is missing both electrons =aparticle =4He++ 4He

  23. Isotopes • Dalton’s theory is modified as a result of Chadwick’s studies of the nucleus and the discovery of the neutron. • All atoms of the same element contain the same number of protons and electrons, but atoms of a given element may have different numbers of neutrons. Isotopes-

  24. Isotopes of Hydrogen

  25. Isotopes of Carbon

  26. Introduction to the Periodic Table

  27. Periodic Table • Arranged according to increasing atomic number (number of protons) • Horizontal Rows – • Vertical Columns – • This arrangement is based on chemical similarities that exist in the vertical columns (groups). These groups are referred to as • This system of arrangement was 1st proposed by Dmitri Mendeleev in 1869. His first table consisted of 62 known elements. He was able to predict the presence of several elements that had not yet been discovered based on his table.

  28. Periodic Table • The name periodic table refers to the fact that as we increase the atomic numbers, every so often an element occurs with properties similar to those of an earlier (lower atomic number) element.

  29. Chemical Families of the Periodic Table • Metals • Nonmetals • Metalloids • The Alkali Metals • The Alaline Earth Metals • Transition Metals • The Oxygen Family • Halogens • Nobel Gases http://www.papernapkin.com/

  30. History –1920s • Following Rutherford’s planetary model of the atom, it was realized that the attraction between the electrons and the protons should make the atom unstable • Neils Bohr (1922) proposed a model in which the electrons would stably occupy fixed orbits, as long as these orbits had special quantized locations

  31. History- 1920s The Bohr Model of the Atom • Neils Bohr expands Rutherford’s model of the atom allowing the electrons to travel in successively larger distinct orbits around the nucleus. The outer orbits hold more electrons than the inner orbits and the outer orbits determine the atom’s chemical properties.

  32. History –The Bohr Model • In the Bohr model, the electron can change orbits, accompanied by the absorption or emission of a photon of a specific color of light.

  33. History- The Bohr Model • Bohr’s Model was used to explain why the negatively charged electrons did not fall into the positively charged nucleus of the atom. • The electrons were only able to occupy distinct energy levels or orbits. “Quantized” energy levels.

  34. History- The Bohr Model • If an atom absorb a specific amount of energy (quantum), the outer shell electrons (valence electrons) could be excited into higher energy states. This excited state is unstable, so the electron releases a photon of energy (quantized) as light. Light of specific wavelength correspond to the energy emitted by the electron dropping back to the ground state.

  35. History- The Bohr Model http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/linesp16.swf

  36. Ions • Ions – • Neutral atoms become ions through the addition or removal of electrons. Atoms form ions to gain stability in their valence electron levels.

  37. Ions • Cations = • Metals want to lose electrons to gain a stable octet • Anions = • Nonmetals want to gain electrons to gain a stable octet

  38. Ions • The resulting charge on the ion is determined based on the number of electrons gained or lost.

  39. Ion Charges and the Periodic Table • The position of an element in the periodic table can help determine the resulting charge on its ion.

  40. Like charges repel Unlike charges attract Electrostatics

  41. Compounds That Contain Ions • Substances that want to lose electrons (metals) will react with substances that want to gain electrons (nonmetals) to forms ions. The resulting ions are attracted to each other through electrostatic forces. 2Na + Cl2 2Na+ + 2Cl-  2NaCl 2 Sodium atoms transfers an electron each to 2 chlorine atoms. The resulting ions (sodium and chloride) are attracted to each other through electrostatic forces.

  42. Ionic Bond = Electrostatic Force

  43. Properties of Ionically Bonded Substances • 1. Composed of metals + nonmetals • 2. High melting point. NaCl mpt = 8000C • 3. Good conductors in liquid or aqueous state. Do not conduct as solids.

  44. Ionic Compounds • A chemical compound must have a net charge of zero if it is neutral! • 1. There must be both positive ions (cations) and negative ions (anions) present. • The numbers of cations and anions must be such that the net charge is zero. NaCl Na+ Cl-

  45. Always check the substances location in the Periodic Table! MgCl2 Mg2+ Cl- Cl- Li3N Li+ Li+ N3- Li+

  46. Nomenclature = Naming Compounds • Binary Compounds- compounds that contain two elements

  47. Binary Ionic Compounds • 1. The cation is always named first and the anion second. • 2. A simple cation (obtained from a single atom) takes its name from the name of the element. For example, Na+ is called sodium in the names of compounds containing this ion. • 3. A simple anion (obtained from a single atom) is named by taking the first part of the element name (the root) and adding –ide. Thus the Cl- ion is called chloride.

  48. Eamples- Binary Ionic Compounds • NaCl  • KI  • CaS  • CsBr • MgO 

  49. Binary Ionic Compounds Containing Transition Metals • Many metals can form more than one type of cation. • Lead (Pb) can form Pb2+ or Pb4+ • Gold (Au) can form Au+ or Au3+ • Iron (Fe) can form Fe2+ or Fe3+ Chemists use Roman numerals to specify the charge on the cation. • Fe2+ = Iron II Fe3+ = Iron III

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