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Add metal hydride naming

Add metal hydride naming. Chapter 2 Atoms, Molecules, and Ions. Atomic Theory of Matter. The theory that atoms are the fundamental building blocks of matter reemerged in the early 19th century, championed by John Dalton. Dalton’s Postulates.

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Add metal hydride naming

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  1. Add metal hydride naming

  2. Chapter 2Atoms, Molecules,and Ions

  3. Atomic Theory of Matter The theory that atoms are the fundamental building blocks of matter reemerged in the early 19th century, championed by John Dalton.

  4. Dalton’s Postulates 1. Each element is composed of extremely small particles called atoms.

  5. Dalton’s Postulates 2. All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements. * He was a little wrong on this one…isotopes!

  6. Dalton’s Postulates 3. Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions. Law of Conservation of Mass

  7. Dalton’s Postulates 4. Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms. Law of Definite Proportions (a.k.a. Law of Constant Composition) See Chapter 1 Notes

  8. Dalton also deduced theLaw of Multiple Proportions CO2 CO *This is always a whole number!

  9. The Electron • J. J. Thompson is credited with their discovery (1897). • Streams of negatively charged particles were found to emanate from cathode tubes.

  10. Millikan Oil Drop Experiment Determined the charge & mass of the electron

  11. Discovery of the Nucleus Ernest Rutherford shot  particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.

  12. The Nuclear Atom • Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom. • Most of the volume of the atom is empty space.

  13. Other Subatomic Particles • Protons were discovered by Rutherford in 1919. • Neutrons were discovered by James Chadwick in 1932.

  14. Subatomic Particles (The mass of an electron is so small we ignore it.)

  15. Symbols of Elements Elements are symbolized by one or two letters.

  16. Atomic Number All atoms of the same element have the same number of protons

  17. Mass Number The mass number is the total number of protons and neutrons in the atom.

  18. Atoms of the same element with different masses. Isotopes have different numbers of neutrons. 11 6 12 6 13 6 14 6 C C C C Isotopes:

  19. Atomic Mass (Weight) of an Atom The actual mass can be calculated for any atom: • Actual Mass (g) = (Mass#)(1.67x10-24g) • This can be converted to an atomic mass by relating the mass to carbon-12 Mass of one atom in amu = Mass of a mole of atoms in grams! *Scientists defined the mole so this would work!

  20. Atomic Mass (Weight) of an Element • A weighted average of atomic masses of all the isotopes of an element.

  21. Use the following data to calculate the atomic mass for the element Magnesium Isotope Atomic Mass of IsotopeAbundance Mg - 24 23.982628 µ 78.600 % Mg - 25 24.963745 µ 10.11 % Mg - 26 25.960802 µ 11.29 % (.78600) (23.982628 g) + (.1011) (24.963745 g) + (.1129) (25.960802 g) 18.850 g + 2.524 g + 2.931 g 24.305 g/mol

  22. Add mass spec here(see end of ppt)

  23. Periodic Table: • A systematic catalog of elements. • Elements are arranged in order of atomic number.

  24. Periodic Table • The rows on the periodic chart are periods. • Columns are groups. • Elements in the same group have similar chemical properties.

  25. Groups These five groups are known by their names.

  26. Periodic Table Nonmetals are on the right side of the periodic table (with the exception of H).

  27. Periodic Table Metalloids border the stair-step line (with the exception of Al and Po).

  28. Periodic Table Metals are on the left side of the chart.

  29. Chemical Formulas The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound.

  30. Types of Formulas • Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound. • Molecular formulas give the exact number of atoms of each element in a compound. • Molecular Formula: C6H12O6 • Empirical Formula: CH2O

  31. Ions • When atoms lose or gain electrons, they become ions. • Cations are positive • Anions are negative

  32. Ionic Bonds Ionic compounds (such as NaCl) are generally formed between metals and nonmetals. *See handout on naming and writing formulas!

  33. Molecular Compounds Molecular compounds are composed of molecules and almost always contain only nonmetals. *See handout on naming and writing formulas!

  34. Diatomic Molecules These seven elements occur naturally as molecules containing two atoms. HONClBrIF

  35. Acids • Compounds containing hydrogen with something that looks like a negative ion (for now at least…..) *See handout on naming and writing formulas!

  36. Organic Molecules (Alkanes) • Organic molecules contain carbon. • Alkanes are the most simple organic compounds. They have chains of carbon with only single bonds surrounded by hydrogen. • They are named by the number of carbons in the chain + the suffix -ane

  37. Alkanes

  38. Alcohol • Other classifications of organic compounds have multiple bonds or other atoms or groups of atoms replacing a hydrogen. • Alcohol has an –OH group instead of a –H off one of the carbons in an alkane

  39. They are named with the same root as the alkanes + the suffix -ol • Starting with propanol they include a number in front of the name indicating which carbon the –OH is bonded to

  40. Mass Spectroscopy

  41. Atomic Mass (Weight) of an Element • A weighted average of atomic masses of all the isotopes of an element.

  42. Use the following data to calculate the atomic mass for the element Magnesium Isotope Atomic Mass of IsotopeAbundance Mg - 24 23.982628 µ 78.600 % Mg - 25 24.963745 µ 10.11 % Mg - 26 25.960802 µ 11.29 % (.78600) (23.982628 g) + (.1011) (24.963745 g) + (.1129) (25.960802 g) 18.850 g + 2.524 g + 2.931 g 24.305 g/mol

  43. Where does that data come from? • Mass Spectroscopy • Atoms that go into the mass spectrometer are organized by mass and the relative amounts are measured. • This is how we found isotopes and disproved part (the only part so far) of Dalton’s Atomic Theory.

  44. Mass Spec for Molybdenium http://www.chemguide.co.uk/analysis/masspec/howitworks.html#top - we will assume the charge is +1…so this is the atomic mass

  45. Estimating Atomic Mass of an Element from its Mass Spectrum 100 23 This scale is usually made by setting the tallest line at 100 All we really care about is this ratio… if we have 123 atoms, 23 would be B-10 and 100 would be B-11

  46. ) • You can either find the % abundance and proceed as you learned 1st year. Or just use the ratios as they are…

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