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Chapter 3: The Atom

Chapter 3: The Atom. Chemistry. Atoms:

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Chapter 3: The Atom

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  1. Chapter 3: The Atom Chemistry

  2. Atoms: • The concept of an invisible, basic particle of matter began with the Greeks as early as 400 BC. When the Greeks believed that all matter could be broken down until a basic particle of matter was reached. These particles could not be subdivided. Democritus called them atoms which means indivisible.

  3. Atomic Theory: • John Dalton developed the atomic theory around 1803. It stands today with only a few minor modifications. • Atom- the smallest unit of an element that can exist either alone or in combination with other atoms like it (element) or different from it (compound).

  4. Law of Conservation of Mass: • Def": Matter is neither created nor destroyed. • When atoms are combined, the mass each supplies is converted to product. • Example: A = 4g; B = 9gA+B=AB AB=13g

  5. Law of Definite and Multiple Proportions • Matter is anything that has mass and takes up space. • Massreactants = Mass products

  6. Law of Definite Proportions • regardless of the amount, a pure compound contains the same elements in the same proportion by mass. • The Law of Conservation of Mass is applied to compounds – mass of the COMPOUND is EQUAL to the SUM of the MASSES of the elements that make up the compound. • H2O molecule • 2 H atoms 1 gram mass ea. 2 grams H H=O • 1 O atom 16 grams mass 16 grams O 1=8

  7. Law of Multiple Proportions • If two elements Form more than one compound Then the ratios of the masses Of the second element which combine With a fixed mass of the first element Will be a ratio of small whole numbers.

  8. Law of Multiple Proportions • If two elements A B Form more than one compound Then the ratios of the masses Of the second element which combine With a fixed mass of the first element Will be a ratio of small whole numbers.

  9. Law of Multiple Proportions • If two elements A B Form more than one compound Then the ratios of the masses Of the second element which combine With a fixed mass of the first element Will be a ratio of small whole numbers.

  10. Law of Multiple Proportions • If two elements A B Form more than one compound AB A B2 etc… Then the ratios of the masses Of the second element which combine With a fixed mass of the first element Will be a ratio of small whole numbers.

  11. Law of Multiple Proportions • If two elements A B Form more than one compound AB A B2 etc… Then the ratios of the masses mass ratio Of the second element which combine With a fixed mass of the first element Will be a ratio of small whole numbers.

  12. Law of Multiple Proportions • If two elements A B Form more than one compound AB A B2 etc… Then the ratios of the masses mass ratio Of the second element which combineB With a fixed mass of the first element Will be a ratio of small whole numbers.

  13. Law of Multiple Proportions • If two elements A B Form more than one compound AB A B2 etc… Then the ratios of the masses mass ratio Of the second element which combineB With a fixed mass of the first elementA Will be a ratio of small whole numbers.

  14. Law of Multiple Proportions • If two elements A B Form more than one compound AB A B2 etc. Then the ratios of the masses mass ratio Of the second element which combineB With a fixed mass of the first elementA Will be a ratio of small whole numbers. (1,2,3….)

  15. C:O and CO2 O:C O:C ratio ratio 1:1 2:1 16 g: 12g 32g:12g

  16. H2O and H2O2 H:O H:O Ratio ratio 2:1 2:2 2g: 16g 2g : 32g

  17. Basic assumptions of the Atomic Theory: • All matter is made up of very small particles called atoms • Atoms of the same element have identical size, mass, and properties. Atoms of different elements have different size, mass, and properties. • Atoms cannot be subdivided, created or destroyed. • Atoms of different elements can combine in simple, whole-number ratios to form compounds. • In chemical reactions, atoms are combined, separated, or rearranged.

  18. Atomic structure: • Nucleus: central part of the atom discovered by Rutherford. • Protons- a positively charged particle with a mass of 1.673 x 10-24g. • Neutrons- an electrically neutral particle with a mass of 1.675 x 10-24 g.

  19. Electrons- located in the region around the nucleus called the electron cloud. They were discovered by Thomson using a Cathode Ray Tube. Crooke saw evidence of them in a CRT shadow, but didn't know what they were. Electrons- negatively charged particles with a mass of 9.110 x 10-28 g.

  20. Inferences on structure: • 1. Because atoms are neutral and electrons are negative, the nucleus must contain a positive charge. • 2. Because electrons are so small compared to the mass of the atom, most of the mass must be in the nucleus.

  21. Nuclear forces: 1. There are attractive forces between the particles in the nucleus which holds the nucleus together. • Proton- proton attraction • Neutron-neutron attraction • Proton-neutron attraction 2. This appears to defy the natural law that like charges repel. Atoms have their own set of natural laws. 3. These are short-range forces that hold the nuclear particles together.

  22. Counting atoms: A. Atomic number and Mass number: • Atomic Number- the # protons in the nucleus. • Mass Number- the # of protons plus neutrons. • Mass number - atomic number +# neutrons. • Mass number is the atomic mass to the nearest whole number. • Example: Silver- AN= 47, MN= 108, Neutrons =61.

  23. Isotopes: 1. Def": Atoms of the same element that have different masses. The number of neutrons changes. 2. Example: Isotopes of hydrogen: • Protium-1 p+ 1 e- (99.985% of hydrogen atoms on earth) • Deuterium-1 p+, 1 e-, 1 n0 (0.015%) • Tritium- 1 p+, 1 e-, 2 n0 radioactive

  24. Hydrogen Isotopes

  25. Although isotopes have different masses, they do not differ significantly in their chemical behavior.

  26. Designating Isotopes • Isotopes (called Nuclides) are specified in 2 ways: • Hyphen Notation -Mass number is written with a hyphen after the name of the element (Ex. Tritium is hydrogen – 3) • Nuclear Symbol – show the composition of the nucleus in a symbol: (atomic #) 92235 (mass#) U

  27. Relative atomic mass: • Because atoms are so small, scientists chose an atom of carbon, arbitrarily, as the standard. • The basis for calculating the atomic mass for all the elements is Carbon-12 is exactly 12 amu. • All other elements are given masses relative to that number. • 1 amu = 1/12 the mass of a carbon-12 atom. (1.660 540 x 10-27 kg)

  28. Average atomic mass: • Since most elements have naturally occurring isotopes, the atomic mass for an element is an average for these mixtures based on the percentage of the isotope’s relative abundance. • How to calculate an average weighted mass: Example: Silver – 2 isotopes. • Ag – 107: 51.82%, AM = 106.9041 amu • Ag – 109: 48.18%, AM = 108.9047 amu • Calculations: • (.5182 x 106.9041) + (.4818 x 108.9047) = 107.8682

  29. Example: Iron – 4 isotopes. • Fe – 54: 5.82%, AM = 53.9396 amu • Fe – 56: 91.66%, AM = 55.9349 amu • Fe – 57: 2.19%, AM = 56.9354 amu • Fe – 58: 0.33% AM = 57.9333 amu • Calculations: • (.0582 x 53.9396) + (.9166 x 55.9349) + • (.0219 x 56.9354) + (.0033 x 57.9333) = • 55.84

  30. Relating mass to number of atoms: • The Mole (mol): • The Mole is the SI unit for the amount of a substance. • It is the amount of substance that contains as many particles as there are atoms in exactly 12 g of Carbon – 12. • Def.: The amount of substance that contains 6.022 x 1023 particles. Exactly 12 grams of Carbon contains 6.022 1367 x 1023 atoms. • Avogadro’s Number: The number of particles in one mole of a pure substance. 6.022 x 1023.

  31. Molar Mass: the mass, in grams, of one mole of a pure substance. • This is also known as the atomic mass and is listed on the periodic table. • Usually written in units of g/mol

  32. Conversions

  33. Conversions a. Moles to Mass How many grams in 4.00 moles of Cadmium? = 449.6 g Or 450. g

  34. Conversions • Grams to moles: • How many moles of Ca are in 5.00g Ca? = 0.125 molCa

  35. Conversions • Moles to atoms: • How many atoms in 2.32 mol Zn? • = 1.40 x1024atoms Zn

  36. Conversions • Atoms to grams: (2 steps) • How many grams of Al do we have if we have 4.02 x 1028 atoms of Al? =6.68 x 104mol =1.80 x 106 g Al

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