Fundamental Chemistry Laws

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Fundamental Chemistry Laws
Law of Conservation of Mass Mass is neither
created nor destroyed
Law of Definite Proportions A given compound
always contains exactly the same proportion of
elements by mass
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
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Law of Definite and Multiple Proportions

• Formula for water is H2O
• If water is decomposed (shown here), then there
  will always be twice as much hydrogen gas (H2)
  formed as oxygen gas (O2).

2 H2O (l) ? O2 (g) 2H2 (g)
Conservation of mass no mass created or
destroyed
3
Daltons Atomic Theory (Early 1800s)

• Each element is made up of tiny particles called
  atoms
• Atoms of a given element are identical atoms of
  different elements different

4
Periodic Table periods rows and groups
columns
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Daltons Atomic Theory

• Each element is made up of tiny particles called
  atoms
• Atoms of a given element are identical atoms of
  different elements different
• Chemical compounds are formed when atoms of
  different elements combine with each other. A
  compound always has the same relative numbers and
  types of atoms (Law of Definite Proportions)
• Chemical reactions involve reorganization of
  atoms not changes in the atoms themselves

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Discovery of Electrons
Deflection of Cathode Rays by an Applied Electric
Field J. J. Thomson, 1897
Millikan oil-drop experiment (1909) determined
charge and mass of electron
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The Plum Pudding Model of the Atom (J.J. Thomson,
early 1900s)
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Expected and Actual Results of Rutherfords
Experiment
Plum-pudding model
New Rutherford model
Eventually, positive particles (protons) were
discovered by Rutherford in 1919 and neutral
particles (neutrons) were discovered by James
Chadwick in 1932
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Modern Atomic Structure
10-12 cm or 10-4 Å
10-8 cm 1-5 Å
Nucleus Protons and Neutrons (most of the
mass) Pink Cloud Electrons (most of the volume)
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Atomic Mass Unit (amu)

• A way to make the numbers more manageable
• 1 amu 1.6606 ? 10-24 g or 6.022 ? 1023 amu 1
  g (Conversion Factor)
• Mass of a proton is 1.0073 amu, 1 amu
• 1.67 ?10-24 g ? 1 amu 1.0087 amu
• 1.6606 ? 10-24 g
• Mass of a neutron is 1.0087 amu, 1 amu
• 1.67 ?10-24 g ? 1 amu 1.0087 amu
• 1.6606 ? 10-24 g
• Mass of an electron is 0.00055 amu, 0 amu
• 9.10939?10-28 g ? 1 amu 0.00055
  amu
• 1.6606 ? 10-24 g

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Atomic Symbols
Mass number
A Z
X
Element symbol
Atomic number

• Atomic number (Z) number of protons
• whole number above symbol on periodic table,
  always the same for a given element (identity of
  element)

• Mass number (A) number of protons number of
  neutrons

• Element Symbol (X) is from periodic table

• In a neutral atom
• the number of protons () the number of
  electrons (-)

• Isotope
• (sometimes represented as X-, ie. F-19 or X,
  ie. 19F
• Atoms with the same number of protons but
  different numbers of neutrons (ie. the mass
  number changes)

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Atomic Mass

• Average atomic mass for all naturally occurring
  isotopes of an element
• Written below element on periodic table
• Sum of each isotope atomic mass times the
  abundance of that isotope
• Useful for determining mass of large quantities
  of atoms with mixed isotope numbers

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Bonding
Bonding
a) Even covalent
Electrons may or may not be evenly shared
b) Uneven polar covalent
c) Completely transferred ionic
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Chemical Bonding

• Molecules/Molecular Compounds
• Sharing electrons between atoms
• Molecules formed with covalent bonds
• Molecular formula and empirical formula

• Ionic Compounds
• Electrical attractions between ions of opposite
  charge
• Ions often pair up to make a neutral ionic
  compound (salt)
• Ionic formula

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Typical Traits of Compounds
Covalent/Molecular
Ionic

• all non-metals
• bond overlap of orbitals (sharing of electrons)
• Does NOT break apart in water
• Exist as gas (g), liquid (l), and solid (s) at
  room temp

• metal and non-metal
• bond - electrostatic attraction of ions
• If soluble Breaks apart into ions in water
• Exist as solids at room temp

Acids In water like ionic As pure compound
like covalent
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Ions

• An atom or groups of atoms that has a net
  positive or net negative charge

• Cations
• Positively charged ions
• Each positive charge is 1 less electron (e-) than
  neutral atom

e.g. 24 12Mg 12 protons () and 12 electrons
(-)   Mg 12 protons () and 11 electrons
(-)   Mg2 12 protons () and 10 electrons (-)
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Anions

• Negatively charged ions
• Each negative charge is 1 more e- than neutral
  atom

e.g. 16 8O 8 protons () and 8 electrons
(-)   O- 8 protons () and 9 electrons
(-)   O2- 8 protons () and 10 electrons (-)
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Formation of Sodium Chloride
Ionic Compound Example
2 Na (s) Cl2 (g) ? 2 NaCl
(s)
sodium chlorine sodium chloride salt
Na 10 e- Cl- 18 e-
Na 11 e-
each Cl 17 e-

• ionic bonds (electrical attraction of ions)
• crystalline
• high melting point 801 C
• high boiling point 1413 C

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Predicting Ionic Charges
Atoms often gain or lose electrons to form
stable ions that have the same number of
electrons as a noble gases

• Metals tend to lose electrons to become cations
• Alkali metals loose 1 e-
• ie. Li, Na, K
• Alkaline earth metals loose 2 e-
• ie. Mg2, Ca2, Sr2
• Transition metals and metals in groups 3A-6A
  often have several stable cationic forms so not
  predictable
• ie. Cu and Cu2
• ie. Fe3 and Fe2

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Predicting Ionic Charges

• Non-metals tend to gain electrons to become
  anions
• Halogens gain 1 e-
• ie. F-, Cl-, Br-
• Group 6A nonmetals gain 2 e-
• ie. O2-, S2-, Se2-
• Group 5A nonmetals gain 3 e-
• ie. N3- and P3-

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Making Ionic Compounds

• Pair up positive and negative ions to make the
  compound neutral with the smallest whole number
  combination of each ion
• Use subscripts to indicate more than one ion

ie. Na and Cl-
NaCl
ie. Fe2 and 2 Cl-
FeCl2
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Breaking Ionic Compounds Into Ions

• Look for one ion in a formula that has mainly one
  stable ion
• ie. Mn2S3, S2- is likely anion
• Multiply the charge of the ion times the
  subscript (this gives you the total amount of
  that charge)
• ie. 2 ? 3 -6 (total negative charge)
• The opposite charge must have the same amount to
  be neutral
• ie. 6 (total positive charge)
• Divide the total opposite charge by the subscript
  on the questionable ion to get the charge on that
  ion
• ie. 6 ? 2 3, therefore Mn3

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Naming Ionic Compounds

• Name cation then anion
• Alkali and Alkaline Earth cations are named with
  the element name
• ie. Na is sodium cation
• ie. Sr2 is strontium cation
• Other Metal cations are named with the element
  name followed by a roman numeral indicating
  charge
• Anions are named by taking the root of the
  element name and adding ide
• ie. O2- is oxide anion
• ie. Cl1- is chloride anion

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Naming Ionic Compounds
Combine cation name and the anion name with a
space
ie. NaCl
is sodium chloride
ie. SrO
is strontium oxide
ie. CuO
is copper(II)oxide
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Memorizing Names of Elements

• Alkali metals
• Alkaline earth metals
• Halogens
• noble gases
• Nonmetals B, C, N, O, F, Si, P, S, Se, Te
• Metals First transition row plus Ag, Au, Hg, Pb

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Polyatomic Ions

• Groups of atoms covalently bonded together that
  have an overall charge
• Memorize common polyatomic ions (names, formuals,
  and charges) Table 2.5 p 59
• Whether the polyatomic ion is the cation or the
  anion, use the name give in Table 2.5
• Parentheses are used if more than one polyatomic
  ion is needed

ie. Na3PO4
sodium phosphate
3 Na and PO43-
ie. (NH4)2SO4
2 NH4 and 1 SO42-
ammonium sulfate
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Acids

• Cation is H and anion is halogen/polyatomic ion
• Made from anion of element
• Name hydro __(root)__ic acid
• ie. HCl hydrochloric acid
• Made from polyatomc ion
• Name __(root)__ ic acid (for ions ending in
  ate)
• Name __(root)__ous acid (for ions ending in
  ite)
• ie. H2SO4 (from SO42-, sulfate ion) sulfuric
  acid
• ie. HNO2 (from NO2-, nitrite ion) nitrous acid

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Naming Binary Molecular Compounds

• Name first element followed by space
• Usually farthest to the left on periodic table is
  first
• If same family, lower one is first
• Write root of second element and add ide
• Use Greek prefixes to indicate how many of each
  element
• mono- one (use only on second element)
• di- two, tri-three, tetra-four, penta-five,
  hexa-six, hepta-seven, octa-eight, nona-nine,
  deca-ten

N2O nitrogen monoxide SF6
sulfur hexafluoride
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Naming Simple Organic Compounds (Carbon Based)

• Root name indicates number of carbon atoms
• meth- one carbon
• eth- two carbons
• prop- three carbons
• but- four carbons
• pent- five carbons
• hex-six carbons
• Suffix ending indicates C-C bonding type and/or
  other atoms besides C and H
• Look at more examples in chapter 8 9

ie. propane contains 3 carbons ie. pentanol
contains 5 carbons
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Acids and Organic Compounds have separate rules!!!