1 Atomic Theory and what we understand today about… Atoms, Molecules, and Ions.
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Transcript of PowerPoint to accompany Chapter 1: Part 2 Atomic theory Introduction: Matter, Measurement and...
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PowerPoint to accompany
Chapter 1:Part 2
Atomic theoryIntroduction:
Matter, Measurement and
Molecules
![Page 2: PowerPoint to accompany Chapter 1: Part 2 Atomic theory Introduction: Matter, Measurement and Molecules.](https://reader037.fdocuments.us/reader037/viewer/2022110209/56649e2d5503460f94b1c380/html5/thumbnails/2.jpg)
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Atomic Theory
The theory that atoms are the fundamental building blocks of matter came into being during the period 1803 to 1807 in the work of an English schoolteacher, John Dalton.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Dalton’s Postulates Each element is composed of extremely small particles called atoms.
All atoms of a given element are identical to one another in mass and other properties, but the atoms of a particular element are different from the atoms of all other elements.
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. This is the basis of the law of conservation of mass (or law of conservation of matter) which states that the total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place.
Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms. This is the basis of the law of constant composition (or law of definite proportions) which states that the relative numbers and kinds of atoms are constant, i.e. the elemental composition of a pure substance never varies.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
The Law of Multiple Proportions Was deduced by Dalton from the preceding four
postulates and states that:
If two elements A and B combine to form more than one compound, the masses of B that can combine with a given mass of A are in the ratio of small whole numbers.
Examples H2O consists of 2 hydrogens to 1 oxygen
H2O2 consists of 1 hydrogen to 1 oxygen
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Cathode Rays and Electrons
Streams of negatively charged particles were found to emanate from cathode tubes.
J. J. Thompson is credited with its discovery in 1897. He determined the charge/mass ratio of the electron
to be 1.78 x 108 Cg-1.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Cathode Ray Tubes (TV’s!)
Cathode- Anode+electons
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Millikan Oil Drop Experiment
Once the charge/mass ratio of the electron was known, determination of either the charge or the mass of an electron would yield the other.
Figure 1.19
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Millikan Oil Drop Experiment
In 1909, Robert Millikan at the University of Chicago determined the charge on the electron.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Radioactivity
The spontaneous emission of radiation by an atom was first observed by Henri Becquerel. It was also studied by Marie and Pierre Curie.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
RadioactivityThree types of radiation were discovered by Ernest Rutherford
particles particles rays
Figure 1.21
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Discovery of the Nucleus
Ernest Rutherford shot particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
The Nuclear Atom
Some particles were deflected at large angles. This led Rutherford to postulate that the atom had a nucleus.
Figure 1.22
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
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.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Subatomic Particles
Protons and electrons are the only particles that have a charge.
Protons and neutrons have essentially the same mass.
The mass of an electron is so small we ignore it.
Table 1.5
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Atomic Numbers, Mass Numbers and Isotopes
The number of protons in the nucleus of an atom of any particular element is called that element’s atomic number (Z).
The mass number (A) of an atom is the total number of protons plus neutrons in that atom.
Mass Number (A)
Atomic Number (Z)
6C12 Symbol of element
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Isotopes
Atoms with identical atomic numbers (Z) but different mass numbers (A), or atoms with the same number of protons which differ only in the number of neutrons are called isotopes.
Examples:
116C
126C
136C
146C
carbon-12 isotope
carbon-14 isotope
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Which Isotopes make up carbon in diamonds, graphite, petroleum and coal?
Where does 14C originate and why is there so little in nature except for living things?
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Atomic Mass
Atomic and molecular masses can be measured with great accuracy with a mass spectrometer.
Figure 1.23
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Average Atomic Mass(commonly called Atomic Mass) We use average masses in calculations,
because we use large amounts of atoms and molecules in the real world.
Average atomic mass is calculated from the fractional abundance of each isotope and mass of that isotope.
For example, the average atomic mass of C - made up mostly of 12C (98.93%) and 13C (1.07%) - is 12.01 u.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Mass Spectrum of Chlorine Atoms
How many protonsdoes Chlorine have?
a) What makes the restof the masses?
b) Why are there
2 different isotopes?
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Mass Spectrum of Chlorine Atoms
What would themass spectrum
of ChlorineMolecules look
like?a) What are the masses? b) Which is least common?
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Periodicity
When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.
Figure 1.25
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Periodic Table A systematic catalogue of elements. Elements are arranged in order of atomic number.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Aluminum wrench, Copper pipe, Lead shot, Gold nuggets
Bromine L+V vial, Iodine crystals, Carbon black + Diamond+ Graphite pencil
Sulfur
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Periodic Table
The rows are called periods. Elements in each row are similar sizes.
The columns are called groups. Elements in columns bond alike and Have similar chemical properties.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Groups
The above five groups are known by their names.
Table 1.7
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Periodic Table
Nonmetals are on the right side of the periodic table (with the exception of H).
Metalloids border the stair-step line (with the exception of Al and Po).
Metals are on the left side of the chart.
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Molecules and Chemical Formulae
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.
Notice how the composition of each compound is given by its chemical formula.
Figure 1.29
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Diatomic Molecules
These seven elements occur naturally as molecules containing two atoms. Which are
gases?
Figure 1.28
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Molecular Compounds
Molecular compounds are composed of molecules and almost always contain only nonmetals.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Types of Formulae
Empirical formulae give the lowest whole-number ratio of atoms of each element in a compound, e.g. HO.
Molecular formulae give the exact number of atoms of each element in a compound, e.g. H2O2.
Structural formulae show which atoms are attached to which within the molecule, e.g. H-O-O-H.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Picturing Molecules
Different representations of the methane (CH4) molecule.
Figure 1.30
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Common Ionic Charges
When atoms lose or gain electrons, they become ions. Cations are positive and are formed by elements
on the left side of the periodic chart. Anions are negative and are formed by elements
on the right side of the periodic chart.
Figure 1.31
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Ionic Compounds
Ionic compounds (such as NaCl) are generally formed between metals and nonmetals.
Figure 1.32
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Writing Ionic Formulae
Because compounds are electrically neutral, one can determine the formula of a compound by: writing the value of the charge on the cation
as the subscript on the anion. writing the value of the charge on the anion
as the subscript on the cation.
Note: if the subscripts are not in the lowest whole number ratio, simplify it, e.g. Ca2O2 would become CaO.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Chemical NomenclaturePositive Ions (Cations)a) Cations formed from metal atoms have the same
name as the metal, e.g. Na+ is the sodium ion.
b) If a metal can form different cations, the positive charge is indicated by a Roman numeral in parentheses following the name of the metal, e.g. Au+ is the gold(I) ion and Au3+ is the gold(III) ion.
c) Cations formed from nonmetal atoms have names that end in -ium, e.g. NH4
+ is the ammonium ion.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Chemical NomenclatureCommon Cations Table 1.8
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Chemical NomenclatureNegative Ions (Anions)
a) The names of the monatomic anions are formed by replacing the ending of the name of the element with -ide, e.g. O2- is the oxide ion, OH- is the hydroxide ion, N3- is the nitride ion.
b) Polyatomic anions containing oxygen (called oxyanions) have names ending in -ate for the most oxidized form e.g. SO4
2- is the sulfate ion or –ite for the more reduced form, e.g. SO3
2- is the sulfite ion.
c) Anions derived by adding H+ to an oxyanion are named by adding the prefix hydrogen e.g. HCO3
- is the hydrogen carbonate ion (bicarbonate ion) or dihydrogen as in dihydrogen phosphate H2 PO4
-2 .
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Chemical NomenclatureCommon Anions
Table 1.9
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Common Oxy-anion
Names, Formulae & Charges
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
K-Iron-CNsalts:
Left: KFe2+(CN)3
(Potassium ferrocyanide)
Right:KFe3+(CN)4
(Potassium ferricyanide)
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Chemical NomenclatureMore on naming oxyanions
Examples:ClO4
- perchlorate ion (one more O atom than chlorate)
ClO3- chlorate (most common oxidized form)
ClO2- chlorite ion (one less O atom than chlorate)
ClO- hypochlorite ion (one O atom less than chlorite)
Which is the most oxidized ion? Which is most reduced?
Figure 1.34
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Chemical NomenclatureName and Formulae of Acids
1. Acids containing anions whose names end in -ide are named by changing the -ide ending to -ic, adding the prefix hydro- to this anion name, and then following with the word acid.
2. Acids containing anions whose names end in -ate or -ite are named by changing the -ate ending to -ic and the -ite ending to -ous and then adding the word acid.
Figure 1.36
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Chemical NomenclatureIonic Compounds
Names of ionic compounds consist of the cation followed by the anion name, e.g. Cu(ClO4)2 is copper(II) perchlorate, and CaCO3 is calcium carbonate.
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Chemical NomenclatureBinary Molecular Compounds
1. The name of the (more reduced) element farther to the left in the periodic table is written first (usually a metal), eg. MnO2 , NaCl, FeFe2 S4 , K(MnO4) .
2. If both elements are in the same group in the periodic table, the one having the higher atomic number (more reduced) is written first eg. SO2 , SiC .
3. The name of the second element is given an -ide ( or –ite or –ate) ending.
4. Greek prefixes are used to indicate the number of atoms of each element.
ExamplesN2O4 is dinitrogen tetroxideP4S10 is tetraphosphorus decasulfide.
Table 1.10
It’s all Greek prefixes to me!
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia