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Dalton’s Atomic Theory

Dalton’s Atomic Theory

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Dalton’s Atomic Theory

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  1. Dalton’s Atomic Theory 1. Elements are made of tiny particles called atoms. 2. All atoms of a given element are identical (not exactly; isotopes) 3. The atoms of a given element are different from those of any other element. 4. 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. 5. 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.

  2. Laws of Chemical Combination Law of Definite Proportions: Different samples of the same compound always contain its constituent elements in the same proportions by mass • This is also called the Law of Constant Composition • Suppose we analyze samples of water from different sources • We will find in each sample the same ratio by mass of hydrogen to oxygen.

  3. So… • A 10.0 gram sample of water, H2O, is comprised of 8.9 grams oxygen and 1.1 grams hydrogen. • Thus, a 100.0 gram sample of water, is comprised of 89.0 grams of oxygen and 11.0 grams of hydrogen.

  4. Laws of Chemical Combination Law of Multiple Proportions (or, Dalton’s Law): If two elements can combine to form more than one compound, then the masses of one element that combine with a fixed mass of the other element are in the ratios of small whole numbers • For instance, the atom carbon forms two stable compounds with the atom oxygen: • Carbon Monoxide (C1O1) and Carbon Dioxide (C1O2) • Neither C1.2O1.3 nor C1.2O2.2, respectively

  5. Formulas of Compounds Chemical Formula: An expression showing the chemical composition of a compound in terms of the symbols for the atoms of the elements involved. Rules for Writing Formulas 1. Each atom present is represented by its element symbol. 2. Metals are written first, followed by non-metals. 3. The number of each type of atom is indicated by a subscript written to the right of the element symbol. 4. When only one atom of a given type is present, the subscript 1 is not written.

  6. Formulas of Compounds Write the formula for each of the following compounds: (a) A molecule contains four phosphorous atoms and ten oxygen atoms. (b) A compound contains one uranium atom and six fluorine atoms. (c) A compound contains one aluminum atom and three chlorine atoms.

  7. The Structure of the Atom Atom: The basic unit of an element that can enter into chemical combination They are made of even smaller particles called subatomic particles (Dalton was wrong!) Electrons Protons Neutrons

  8. Summary of the Structure of the Atom Mass and Charge of Subatomic Particles ParticleMass (grams)Charge Unit Electron 9.1095 x 10-28 -1 Proton 1.67252 x 10-24 +1 Neutron 1.67495 x 10-24 0 • Nucleus in center of atom – houses protons & neutrons • Electrons around nucleus in orbitals

  9. Atomic Symbols Atomic Number: The number of protons in the nucleus of an atom. Mass Number: The total number of neutrons (variable) and protons present in the nucleus of an atom.

  10. Atomic Symbols • We use the symbol to represent the atom • X = the symbol of the element • A = the mass number • Z = the atomic number • How do we calculate the number of neutrons? • Mass Number = • Protons + Neutrons • Thus, neutrons = A - Z

  11. Atomic Symbols: Practice • Give the atomic symbol for a species that houses 38 protons and 48 neutrons. • Give the atomic symbol for a species that houses 38 protons and 50 neutrons. • What is the difference?

  12. Isotopes • Atoms having the same atomic number but different mass numbers; i.e., different number of neutrons. • The Periodic Table houses average mass numbers based on percent abundance of each isotope.

  13. Isotopic abundance • % abundance = (# atoms of isotope/total # of atoms of all isotopes) x 100% • For ex: element X has two isotopes • X-20 and X-21 • Out of every 100 atoms of X, 90 are X-20 and the rest (10) are X-21 • So, X-20 has an abundance of (90/100) x 100% = 90% • And X-21 has an abundance of 10%

  14. Atomic weight • Atomic weight = (% abundance isotope/100)(mass of isotope 1) + … • An average measurement • Shown on the periodic table • Measured in atomic mass units (amu) • But more commonly in molar mass or grams/mole • Will discuss anon

  15. Practice:What’s the atomic weight? • B-10 = 19.91% abundance • 10.0129 amu = mass • B-11 = 80.09% abundance • 11.0093 amu = mass

  16. Atomic mass unit • Amu or u = 1/12th mass of C-12 • = 1.661 x 10-24 g • (C-12= 12 amu) • Used for subatomic particles

  17. The mole • Mole = amt that contains as many “things” as there are atoms of 12 g of C-12 • 1 mole = 6.022 x 1023 particles • Molar mass (MM) = mass in grams per 1 mole of particles (g/mol)

  18. Practice • How much does one atom of argon (Ar) weigh?

  19. Practice • How many moles and atoms of Ar are in a 5.00 gram sample?

  20. Your turn • How many grams of niobium (Nb) are there in a 2.00 mole sample? • How many atoms?

  21. Introduction to the Periodic Table

  22. What do the numbers mean? • Each cell houses: • Elemental symbol • Elemental name • Atomic number • Molar mass • Sundry information • Let’s take a look

  23. Introduction to the Periodic Table • Elements symbolized • by either one capital letter • or two letters; first capital, second lower-case • Examples: • B = boron • Cl = chlorine • P = phosphorus • Na = Sodium • Symbols don’t always match up with English names • Because use Greek and Latin words

  24. Introduction to the Periodic Table • Groups: The elements in a vertical column of the periodic table • “Old” School: Group-A and Group-B • Taller groups: A • Shorter groups: B • “New” School: Groups I - XVIII • Elements classified in groups based on valence electron configuration: similar chemical properties • Group names: • Alkali Metals: The Group 1A elements (LiFr) • Alkaline Earth Metals: The Group 2A elements (BeRa) • Halogens: The Group 7A elements (FAt) • Noble Gases: The Group 8A (HeRn) • Transition Metals: The Group-B elements

  25. Introduction to the Periodic Table • Period: The elements in each horizontal row of the periodic table • There are 7 periods • Increase in atomic number going to the right • Two rows at bottom are: • Lanthanides or rare-earth metals • Actinides (contain transuranium metals)

  26. Introduction to the Periodic Table • Elements can be divided into 3 categories: • Metals • Metalloids • Nonmetals • Step-ladder divides metals from nonmetals • Metalloids surround step-ladder • Periodic Table roughly 75% metals • Let’s label the Periodic Table handout!

  27. Introduction to the Periodic Table: Metals 1. Efficient conduction of heat and electricity 2. Malleability (they can be hammered into thin sheets) 3. Ductility (they can be pulled into wires) 4. A lustrous (shiny) appearance

  28. Introduction to the Periodic Table: Metalloids • Or semimetals, have properties that fall between those of metals and nonmetals

  29. Introduction to the Periodic Table: Nonmetals • Usually poor conductors of heat and electricity • More varied physical properties than metals

  30. Natural States of the Elements • All metals and metalloids are solid • Except mercury = quicksilver • Non-metals both solid and gaseous • Noble gases • H, O, N, F, Cl • Some nonmetals are diatomic • A molecule that consists of two similar atoms • All halogens • Oxygen, nitrogen, and hydrogen • One is liquid • Bromine

  31. Natural States of the Elements Allotropes: Two or more forms of the same element that differ significantly in chemical and physical properties. Example Phosphorus: red and white Carbon: Diamond, Graphite, Buckyball