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Unit II : Atoms, Molecules, Ions and Nuclear Chemistry

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1 Unit II : Atoms, Molecules, Ions and Nuclear Chemistry

2 Early History of Chemistry
Before 16th Century Alchemy: Attempts (scientific or otherwise) to change cheap metals into gold 17th Century Robert Boyle: First “chemist” to perform quantitative experiments 18th Century George Stahl: Phlogiston flows out of a burning material. Joseph Priestley: Discovers oxygen gas, “dephlogisticated air.”

3 Law of Conservation of Mass
Discovered by Antoine Lavoisier Mass is neither created nor destroyed Combustion involves oxygen, not phlogiston

4 Law of Definite Proportions
Sometimes called the law of constant composition John Proust (1799) A given compound always contains exactly the same proportion of elements by mass. Carbon tetrachloride is always 1 atom carbon per 4 atoms chlorine.

5 Law of Multiple Proportions
When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers. The ratio of the masses of oxygen in H2O and H2O2 will be a small whole number (“2”).

6 Dalton’s Atomic Theory
Each element is made up of tiny particles called atoms. The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways. Chemical compounds are formed when atoms combine with each other. A given compound always has the same relative numbers and types of atoms. Chemical reactions involve reorganization of the atoms - changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction.

7 Avagadro’s Hypothesis (1811)
At the same temperature and pressure, equal volumes of different gases contain the same number of particles. 5 liters of oxygen 5 liters of nitrogen SAME NUMBER OF PARTICLES

8 Chemical Symbols Co Cu He H
The chemical symbols used today were developed by Jons Jakob Berzelius. They consist of one or two letters. The first letter is always capitalized. The second, if there is one, is never capitalized. The second letter is often a letter prominent in the pronunciation. Co Cu He H

9 Early Experiments to Characterize the Atom
Democritis Dalton Thomson Rutherford Bohr Schrodinger

10 Cathode Ray Streams of negatively charged particles were found to emanate from cathode tubes. J. J. Thompson is credited with their discovery (1897).

11 Deflection of Cathode Ray by an Applied Electric Field

12 Rutherford Gold Foil Experiment
Ernest Rutherford shot  particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.

13 Rutherford Gold Foil Experiment
Expected results if Thomson’s Model was correct Actual results

14 Summary of Atom Democritis - first idea of the atom
Dalton - ~1807 Atomic Theory (atom = ball) Thomson - Experiment with Cathode Ray Discovered the electron “Plum pudding model” Rutherford - Gold Foil Experiment Atom is mostly empty space with a dense, positively charged nucleus Bohr - “Solar System” Schrodinger - Quantum Mechanical Model

15 Modern View of Atomic Structure
The atom contains 3 types of subatomic particles: Electrons Negatively charged, found outside the nucleus… very small mass… would take ~2000 electrons to equal the mass of 1 of the other subatomic particles Protons Positively charged, found in the nucleus Neutrons Found in the nucleus… like a proton but with no charge

16 Subatomic Particle Summary
Mass (g) Charge Atomic Mass Scale Relative Mass Relative Charge Proton (p+) 1.6726x10-24 +1.602x c amu 1 +1 Neutron (n0) amu Electron (e-) 9.109 x 10-27 -1.602x c 5.486 x amu 1/1836 -1

17 Useful Units Atomic Diameter:10-10m (1 Å) = 100 pm = 1x10-8cm
Nuclear Diameter: 10-13cm 1 amu (atomic mass unit) = x10-24 kg Based off of C-12… relative mass of O w/ 8 protons and 8 neutrons is 1 picometer (pm) = 1 x m NOTE: Heaviest atom has a mass of only 4.8x10-22g and a diameter of only 5x10-10m Biggest atom is 240 amu and is 50 Å across. Typical C-C bond length 154 pm (1.54 Å) Molecular models are 1 Å /inch or about 0.4 Å /cm

18 K Atomic Number A = Z + N 39 19 Element Symbol
A = mass number = number of protons + number of neutrons K 39 Element Symbol 19 Z = atomic number = number of protons A = Z + N Mass number = # p+ + # n0

19 Isotopes Isotopes are atoms of the same element with different masses.
Isotopes have different numbers of neutrons. Carbon has two natural isotopes: 12C and 13C How does this change things?

20 Isotopes On the periodic table the given atomic mass is the AVERAGE mass We use percentage to help with this math. For Carbon, there is 98.89% 12C and 1.11% of 13C Now what? 12 (mass of 12C) x = 13 (mass of 13C) x = =

21 Isotopes On the periodic table the given atomic mass is the AVERAGE mass We use percentage to help with this math. For Carbon, there is 98.89% 12C and 1.11% of 13C

22 Mass Spectrometer Used to experimentally determine mass and percent abundances of isotopes

23 The Periodic Table Noble Gases Alkali Metals Main Group
Alkaline Earths Halogens Transition Metals Lanthanides and Actinides General Chemistry: Chapter 2 Prentice-Hall © 2002

24 Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.

25 The Periodic Table 1869- Dmiti Mendeleev organized 1st periodic table by lining elements up in horizontal rows in order of increasing atomic weight Left spaces when an element was not known but should exist and have properties similar to the element about it in his table. Periodicity – periodic repetition of the properties of elements 1913- H.G.J. Moseley organized elements according to increasing atomic number Law of chemical periodicity The properties of the elements are periodic functions of atomic numbers.

26 Features of the Periodic Table
Vertical Columns: Groups or Families A and B groups. “A” main group, “B” transition elements Horizontal Rows: called periods Periodic Law: when the elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties

27 Metals ~80% of all elements are metals
High electrical conductivity and have a high luster when cleaned Ductile (can be made into wires) and Malleable Can form alloys Are solid at room temperature except for one… which one is it?

28 Nonmetals Generally nonlustrous… which means not shiny
Generally poor conductors of electricity With the exceptions of carbon (graphite), none conduct electricity Allotropes A particular element can often exist in several different forms Carbon as graphite or diamond… oxygen as O2 or O3 (ozone)

29 Metalloids Semi-metals

30 Overview of Groups Group 1A Group 2A Alkali Metals
Metals, solid at room temp, and VERY REACTIVE Only found in nature combined in compounds Group 2A Alkaline Earth Metals Metals, solid at room temp, and also only found in nature in compounds Except for Be, all react with water to produce alkaline solutions

31 Overview of Groups Group 7A Group 8A Halogens
All exist in diatomic molecules Combine violently with alkali metals to form salts Group 8A Noble Gases (Inert Gases) Least reactive elements All are gases

32 Overview of Groups Group B Metals (Groups 3-12):
Transition Metals & Inner Transition Metals Lanthanides- shiny metals similar in reactivity to alkaline earth metals Actinides- unique in nuclear structures – nuclei are unstable and therefore radioactive Inner Transitions Metals – rare-earth metals

33 The Periodic Table Noble Gases Alkali Metals Main Group
Alkaline Earths Halogens Transition Metals Lanthanides and Actinides General Chemistry: Chapter 2 Prentice-Hall © 2002

34 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. Molecular compounds are composed of molecules and almost always contain only nonmetals.

35 Chemical Formulas: Diatomic Molecules
These 7 elements occur naturally as diatomic molecules

36 Types of Formulas Molecular Formula – CH4 Structural Formula
Ball and Stick Model Space-Filling Model

37 Types of Formulas Molecular Formula Empirical Formula
Exact number of each type of atom C6H12O6 Empirical Formula Lowest whole number ratio of each type of atom CH2O


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