History of Atomic Theory
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1 Unit 2 The Atom
2 History of Atomic Theory A. Democritus and Aristotle Democritus named the "atom" - means indivisible
3 Dalton (with work of Lavoisier, Proust, and Gay-Lussac) 1. atomic theory - first based on evidence a. all matter is composed of atoms b. all atoms are alike in a particular element c. atoms combine in simple whole number ratios to form compounds
4 Atomic theory is based on three laws a. law of conservation of mass - matter cannot be created or destroyed b. law of definite proportions - compounds have same mass ratio of elements c. law of multiple proportions - compounds have same whole number ratio of elements Cr 2 O 3 CrO 3
5 John Dalton s Atom Atoms are solid and indivisible.
6 Atomic Particles Thomson - discovered the electron, mass of proton, plum pudding model (particles mixed) Millikan - discovered the electron s charge to mass ratio. Chadwick - discovered the neutron
7 Components of the Atom Electrons J.J. Thomson found in 1897 Particles deflected in Cathode Ray tube Negative one charge Very small mass 1/2000 that of the lightest atom
8 J. J. Thomson s Atom Thompson hypothesized that all atoms had such negative charges within. (plum pudding model) Later, scientists proved that atoms contained both positive and negative charges (which balance each other out).
9 Ernest Rutherford 1911 Gold Foil Experiment using alpha particles Found the nucleus is positively charged and contains most of the atom s mass. Thought the rest of the atom was mostly empty space.
10 Bohr Studied relationships between energy, wavelength, and frequency Used Rutherford s discovery of the nucleus Developed the planetary orbit model of the atom
11 Niels Bohr s Atom Electrons orbit the nucleus in orbits, like a solar system. Planetary Model
12 The Atom Atomic Theory
13 Parts of the atom A. proton p +, neutron n o and electron e - B. Terms 1. atomic number (Z) = number of protons (identity) 2. # electrons = number of protons 3. mass number = # protons plus # neutrons (A) 4. # neutrons = mass # - atomic #
14 Isotopes 1. atoms with the same atomic number (Z), but different mass number (A); they have different number of neutrons 2. mass number is written after the element name, i.e. carbon nuclear symbols put mass number on top, atomic number on bottom out to the left of the chemical symbol Examples: carbon (see board)
15 Atomic Mass 1. based on carbon-12 having 12 atomic mass units (a.m.u.) 2. average atomic mass is a weighted average of all isotopes Steps to calculate: 1. multiply relative abundance and average mass for each isotope (to correct sig figs). 2. add together (to least accurate decimal place) to get average atomic mass for the element.
16 Example: What is the atomic mass of silicon if 92.21% is mass , 4.70% is mass , and 3.09% is mass ? (.9221)( amu) = (.0470)( amu) = 1.36 (.0309)( amu) = Add together for average atomic mass
17 The Atom Using the Mole
18 Scientific Notation Expressed as product of number and power of 10; such as 3.5 x 10 3 cm Removes doubt about significant figures - all numbers are significant. Conversion: Only one nonzero number in front of the decimal. Find the power of power of 10 for each space moved to the left -1 power of 10 for each space moved to the right Follow normal rules about sig figs to determine how many in answer. Use exponent key on calculator (EE or EXP) to REPLACE the x 10 part of number (4.2 x 10 2 cm)(3.0 x 10-5 cm) = Put in: 4.2 EXP 2 x 3.0 EXP -5 (or EE) (or EE)
19 Formula and Molecular Mass Sum of atomic masses of all atoms in a compound Steps to calculate: 1. write the number of atoms of each element present in the formula 2. find mass of each element on table 3. multiply these two together to get total mass of each element in compound 4. add masses together to least accurate decimal place Example: (see board)
20 The Mole A mole is always the same number of items just like a dozen is always twelve but the number is HUGE! Avogadro s number is the number of particles in a mole. It is x This helps us count the number of particles by weighing. The mass of one mole of particles is equal to its atomic or formula mass in grams. So 3 Equivalents: x particles = 1 mol = # g from table (Use as converting ratios put in the units before any numbers! )
21 Mole Examples 1. g to moles: How many moles are there in 50.4 g Ca? 50.4 g Ca 1 mol Ca = 1.26 mol Ca 40.08g Ca 2. particles to moles: How many moles of NaCl are there if there are 5.68 x formula units? 5.68 x formula units 1 mol NaCl x formula units = 9.43 x 10-6 mol NaCl 3. g to particles: How many molecules of H 2 O in 42.3 g of water? 42.3g H 2 O x molecules g H 2 O = 1.41 x molecules H 2 O
22 End of Unit 2
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