Particle Theory of Matter

1
Particle Theory of Matter

• Democritus 400 B.C. first to state that all
  matter is made of small, hard particles called
  atoms comes from the Greek word atomos
  which means indivisible.
• Aristotle did not agree, felt that all matter
  was a flowing continuous material in which
  everything in connected his theory was accepted
  for 2000 years
• Elements Earth, Wind, Fire, Water
• Properties Coldness, Hotness, Dryness, Moistness
• Neither theory was supported by experimental
  evidence merely observation

2
Alchemy/Phlogiston Theory of Fire

• Alchemy preceded modern chemistry
• Began in Egypt, Persia, and Mesopotamia
• Began as a mixture of magic, religion, and
  science
• In the Middle Ages, interested mainly in the
  transmutation of metals turning lead into
  gold
• Alchemy lead to many important advances in
  equipment, procedures, and specifically in the
  areas of metallurgy, dyeing, glass-making and
  medicine
• Phlogiston Theory of Fire During the 1700s, it
  was theorized that a substance called phlogiston
  was contained in flammable substances When a
  fire went out it was smothered with phlogiston
  that was released when something burned.

3
Foundations of Atomic Theory

• In the 1700s, the particle theory of matter was
  becoming more accepted an element was a
  substance that could not be broken down.
• There was controversy over whether compounds were
  made up of the same ratio of elements each time
  they were formed.
• In the 1770s, new technology allowed for better
  quantitative measurements to establish the fact
  that the ratios were the same every time.
• Three new laws were established as a result of
  the new technology the so-called Laws of Matter

4
Law of Conservation of Mass
Mass is neither created nor destroyed during
chemical or physical reactions.
Total mass of reactants Total mass of products
Antoine Lavoisier
5
Laws of Matter

• Law of Conservation of Matter
• When a chemical change (reaction) or physical
  change takes place, matter is neither created nor
  destroyed. (Massreactants Massproducts)
• Law of Definite Proportions
• When two elements react to form a compound, the
  resulting chemical compound contains the same
  proportions of elements by mass regardless of the
  size or source of the compound.
•  
• Law of Multiple Proportions
• When two elements (C and O) can combine to form
  more than one compound (say, CO and CO2), then
  for a fixed mass of C, the masses of O in the two
  different compounds always form a whole number
  ratio.

6
Daltons Atomic Theory (1808)

• All matter is composed of extremely small
  particles called atoms
• Atoms of a given element are identical in size,
  mass, and other properties atoms of different
  elements differ in size, mass, and other
  properties

John Dalton

• Atoms cannot be subdivided, created, or destroyed
• Atoms of different elements combine in simple
  whole-number ratios to form chemical compounds
• In chemical reactions, atoms are combined,
  separated, or rearranged

7
Modern Atomic Theory
Several changes have been made to Daltons theory.
Dalton said
Atoms of a given element are identical in size,
mass, and other properties atoms of different
elements differ in size, mass, and other
properties
Modern theory states
Atoms of an element have a characteristic average
mass which is unique to that element.
8
Isotopes
Elements occur in nature as mixtures of
isotopes. Isotopes are atoms of the same element
that differ in the number of neutrons Isotopes
differ in mass but do not differ significantly in
chemical behavior
9
Modern Atomic Theory 2
Dalton said

• Atoms cannot be subdivided, created, or destroyed

Modern theory states

• Atoms cannot be subdivided, created, or destroyed
  in ordinary chemical reactions. However, these
  changes CAN occur in nuclear reactions!

10
Discovery of the Electron
In 1897, J.J. Thomsen used a cathode ray tube to
deduce the presence of a negatively charged
particle.
Cathode ray tubes pass electricity through a gas
that is contained at a very low pressure.
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Discovery of the Electron (cont.)

• Cathode rays deflected by the magnetic field and
  by the negatively-charged source.
• Proved the presence of negatively-charged
  particles eventually called electrons
• Calculated the ratio of the charge to the mass of
  the electrons it was always the same ratio 11

12
Some ModernCathode Ray Tubes
13
Thomsens Atomic Model
Thomsen believed that the electrons were like
plums embedded in a positively charged pudding,
thus it was called the plum pudding model.
14
Mass of the Electron

• 1909 Robert Millikan determines
• Presence of electrons in an atom
• Mass of the electron
• Neg. charge on an electron
• Atoms are divisible can be broken apart

The oil drop apparatus
Mass of the electron is 9.109 x 10-31 kg 1/1837
the mass of a proton
15
Conclusions from the Study of the Electron

• Cathode rays have identical properties regardless
  of the element used to produce them. All elements
  must contain identically charged electrons.
• Atoms are neutral, so there must be positive
  particles in the atom to balance the negative
  charge of the electrons
• Electrons have so little mass that atoms must
  contain other particles that account for most of
  the mass

16
Rutherfords Gold Foil Experiment

• Alpha particles are helium nuclei (nucleus w/no
  electrons
• Particles were fired at a thin sheet of gold foil
• Particle hits on the detecting screen (film) were
  recorded

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Try it Yourself!
In the following pictures, there is a target
hidden by a cloud. To figure out the shape of the
target, we shot some beams into the cloud and
recorded where the beams came out. Can you figure
out the shape of the target?
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The Answers
Target 2
Target 1
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Rutherfords Findings

• Most of the particles passed right through
• A few particles were deflected
• VERY FEW were greatly deflected

Like howitzer shells bouncing off of tissue
paper!
Conclusions

• The nucleus is small
• The nucleus is dense and massive
• The nucleus is positively-charged

20
Nuclear Structure

• Nucleus is held together by nuclear forces
  between protons, neutrons, and each other when
  these particles are close together they establish
  strong attraction called the strong nuclear
  force.
• Atoms are electrically neutral so the proton is
  equal in number but opposite in charge to the
  electron.
• Scientists could not explain the differences in
  mass so they decided that there must be a neutral
  particle with a mass similar to that of a proton
  neutron.
• James Chadwick discovered the neutron in 1932

21
Atomic Particles
1 amu 1/12 the mass of a Carbon-12 atom
22
The Atomic Scale

• Most of the mass of the atom is in the nucleus
  (protons and neutrons)
• Nuclear density is very high 2 X 108 metric
  tons/cm3
• Electrons are found outside of the nucleus (the
  electron cloud)
• Most of the volume of the atom is empty space
• Atomic radii from 40 to 270 pm (10-12)
• Nuclear radii 0.001 pm

q is a particle called a quark
23
About Quarks
Protons and neutrons are NOT fundamental
particles.
Protons are made of two up quarks and one
down quark.
Neutrons are made of one up quark and two
down quarks.
Quarks are held together by gluons
24
Atomic Number
Atomic number (Z) of an element is the number of
protons in the nucleus of each atom of that
element.
25
Isotopes
Isotopes are atoms of the same element having
different masses due to varying numbers of
neutrons. (Nuclide term used to
describe any isotope of any element)
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Mass Number
Mass number (A) is the number of protons and
neutrons in the nucleus of an isotope.
Mass p n0
8
8
18
18
Arsenic
75
33
75
Phosphorus
15
31
16
Nuclear Symbol A Z
Mass Atomic
X
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Bonds

• Chemical bond mutual chemical attraction
  between the nuclei (protons) and the valence
  electrons of different atoms that binds the atoms
  together

• Ionic bonds bonding that results from the
  electrical attraction of large numbers of cations
  and anions (involves the transfer of electrons)
• Covalent bonds bonding that results from the
  sharing of electron pairs between atoms

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Ions

• Ion Atom that has lost or gained electrons to
  take on a noble gas configuration (stable state)
• Cation A positive ion (lost es)
• Mg2, NH4
• Anion A negative ion (gained es)
• Cl-, SO42-
• Ionic Solid (salt) solids composed of
  oppositely-charged ions (NaCl)
• Polyatomic ion Group of atoms covalently-bonded
  together that have lost or gained es to exhibit
  an overall charge.

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Periodic Table with Group Names
30
Periodic Table

• Periods
• Period horizontal row on the PT
• Valence electrons electrons located in the
  outside energy level
• Tells us the energy level of the outside
  (valence) electrons in an element (Principle
  Quantum )
• Can be broken down into subdivisions that
  correspond to the s, p, d, and f orbitals
• Groups
• S-block, Groups 1 and 2
• Group 1 Alkali Metals
• 1 s electron (valence electron) in outside
  energy level (willing to give up their electron -
  1 ions)
• Soft, silvery metals with low density and low
  mps
• Highly reactive, never found as an element in
  nature, always in compounds (NaCl salt)

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Periodic Table

• Groups (cont.)
• Group 2 Alkaline Earth Metals
• 2 s electrons (valence electrons) in outside
  energy level (willing to give up their electrons
  - 2 ions)
• Denser, harder, stronger, less reactive than
  Group 1
• Abundant in the Earths Crust (found in
  compounds)
• D-block, Groups 3-12 Transition Metals
  (Elements)
• Shiny luster when polished, malleable, ductile,
  typical metallic properties
• High mps and bps, thermally and electrically
  conductive
• Valence electrons in the d-orbital
• Group sum of s and d electrons

32
Periodic Table

• Groups (cont.)
• p-block, Groups 13-18
• Properties vary greatly
• Metals soft, less dense than d-block
  metals/harder and more dense than s-block metals
• Metalloids Brittle solids with some metallic
  and nonmetallic properties semiconductors
• B, Si, Ge, As, Sb, Te
• Nonmetals
• Halogens (Group 17) most reactive non-metals
• Gases F,Cl
• Liquid Br
• Solid I
• Noble Gases (Group 18) inert (unreactive)
  gases full outside energy level

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Periodic Table

• Groups (cont.)
• F-block, Lanthanides and Actinides
• Lanthanides shiny metals similar in reactivity
  to Group 2
• Actinides all radioactive, most are synthetic,
  only Thorium and Uranium are found naturally

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Predicting Ionic Charges
Group 1
Lose 1 electron to form 1 ions
H
Li
Na
K
35
Predicting Ionic Charges
Group 2
Loses 2 electrons to form 2 ions
Be2
Mg2
Ca2
Ba2
Sr2
36
Predicting Ionic Charges
Group 13
Loses 3 electrons to form 3 ions
B3
Al3
Ga3
37
Predicting Ionic Charges
Group 14
Lose 4 electrons or gain 4
electrons?
Neither! Group 13 elements rarely form ions.
38
Predicting Ionic Charges
Nitride
N3-
Group 15
Gains 3 electrons to form 3- ions
P3-
Phosphide
As3-
Arsenide
39
Predicting Ionic Charges
Oxide
O2-
Gains 2 electrons to form 2- ions
Group 16
S2-
Sulfide
Se2-
Selenide
40
Predicting Ionic Charges
F1-
Br1-
Fluoride
Bromide
Group 17
Gains 1 electron to form 1- ions
Cl1-
Chloride
I1-
Iodide
41
Predicting Ionic Charges
Group 18
Stable Noble gases do not form ions!
42
Predicting Ionic Charges
Groups 3 - 12
Many transition elements
have more than one possible oxidation state.
Iron(II) Fe2
Iron(III) Fe3
43
Predicting Ionic Charges
Groups 3 - 12
Some transition elements
have only one possible oxidation state.
Zinc Zn2
Silver Ag
Cadmium Cd2
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Oxidation Numbers
Oxidation s arbitrary number of charge
assigned to indicate the general distribution of
es among atoms in a covalent compound or
polyatomic ion Oxidation Number Rules (in
general) Elements 0 Alkali Metals
1 Alkaline Earth Metals 2 Aluminum,
Boron, Gallium 3 Halogens -1 Oxygen
Group -2 Nitrogen Group -3 Oxidation s
are useful in determining chemical formulas or in
balancing chemical reactions
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Writing Ionic Compound Formulas
Example Barium nitrate
1. Write the formulas for the cation and anion,
including CHARGES!
( )
Ba2
2. Check to see if charges are balanced.
NO3-
2
Not balanced!
3. Balance charges , if necessary, using
subscripts. Use parentheses if you need more than
one of a polyatomic ion.
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Writing Ionic Compound Formulas
Example Ammonium sulfate
1. Write the formulas for the cation and anion,
including CHARGES!
( )
NH4
SO42-
2. Check to see if charges are balanced.
2
3. Balance charges , if necessary, using
subscripts. Use parentheses if you need more than
one of a polyatomic ion.
Not balanced!
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Writing Ionic Compound Formulas
Example Iron(III) chloride
1. Write the formulas for the cation and anion,
including CHARGES!
Fe3
Cl-
2. Check to see if charges are balanced.
3
3. Balance charges , if necessary, using
subscripts. Use parentheses if you need more than
one of a polyatomic ion.
Not balanced!
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Writing Ionic Compound Formulas
Example Aluminum sulfide
1. Write the formulas for the cation and anion,
including CHARGES!
Al3
S2-
2. Check to see if charges are balanced.
2
3
3. Balance charges , if necessary, using
subscripts. Use parentheses if you need more than
one of a polyatomic ion.
Not balanced!
50
Writing Ionic Compound Formulas
Example Magnesium carbonate
1. Write the formulas for the cation and anion,
including CHARGES!
Mg2
CO32-
2. Check to see if charges are balanced.
They are balanced!
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Writing Ionic Compound Formulas
Example Zinc hydroxide
1. Write the formulas for the cation and anion,
including CHARGES!
( )
Zn2
OH-
2. Check to see if charges are balanced.
2
3. Balance charges , if necessary, using
subscripts. Use parentheses if you need more than
one of a polyatomic ion.
Not balanced!
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Writing Ionic Compound Formulas
Example Aluminum phosphate
1. Write the formulas for the cation and anion,
including CHARGES!
Al3
PO43-
2. Check to see if charges are balanced.
They ARE balanced!
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Naming Ionic Compounds

• 1. Cation first, then anion
• 2. Monatomic cation name of the element
• Ca2 calcium ion
• 3. Monatomic anion root -ide
• Cl- chloride
• CaCl2 calcium chloride

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Naming Ionic Compounds(continued)
Metals with multiple oxidation states

• - some metal forms more than one cation
• - use Roman numeral in name
• PbCl2
• Pb2 is cation
• PbCl2 lead(II) chloride

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Naming Binary Covalent Compounds

• - Compounds between two nonmetals
• - First element in the formula is named first
• - Second element is named as if it were an anion
  (-ide ending)
• - Use prefixes (see hand-out)
• - Use mono- on second element only

P2O5
diphosphorus pentoxide
CO2
carbon dioxide
CO
carbon monoxide
N2O
dinitrogen monoxide
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Naming Acids

• Acids are divided into two groups
• Binary - consist of two elements
• Oxyacids - consist of 3 elements, one of which is
  oxygen.

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Naming Acids

• NAMING BINARY ACIDS
• The name of the binary acid consists of two
  words.
• The first word has three parts
• hydro prefix
• root of the nonmetal element
• ic ending
• The second word is always acid
• Examples
• HCl hydro chlor ic acid hydrochloric acid
• HBr hydro brom ic acid hydrobromic acid
• HF hydro fluor ic acid hydrofluoric acid

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Naming Acids

• NAMING OXYACIDS
• These are more difficult to name because these
  acids have hydrogen, a nonmetal, and may have
  varying numbers of oxygen atoms.
• For example, H2SO5, H2SO4, H2SO3, and H2SO2 are
  all acids. How do we name them? To begin, we need
  a point of reference. Our reference point is
  this
• If the ion ends with ate, remove the ending and
  add ic
• If the Ion ends with ite, remove the ending and
  add ous
• If the ion end with ide, remove the ending and
  add ic
• The second word is always acid.
• Examples
• SO42- sulfate ion H2SO4 sulfuric acid
• NO2- nitrite ion HNO3 nitrous acid
• CN- cyanide ion HCN cyanic acid