Fundamental Chemistry Laws

1
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
2
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 formed
  as oxygen gas.

3
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

4
Periodic Table periods rows and groups
columns
5
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
6
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
Nucleus Protons and Neutrons (most of the
mass) Pink Cloud Electrons (most of the volume)
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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

• 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)
• Atoms with the same number of protons but
  different numbers of neutrons (ie. the mass
  number changes)

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Using Atomic Symbols
Determine the number of protons, electrons, and
neutrons in the following atom.
47 22
Ti
1) Atomic number (Z) of protons 22 2) In a
neutral atom (no charge) of protons of
electrons 22 3) Mass number (A) of protons
neutrons 47
22 N 47
N 25 neutrons
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Isotopes

• Atoms with the same number of protons but
  different numbers of neutrons
• ie. the mass number changes
• Represented by showing a different mass number
• ie.
• ie. 47Ti and 48Ti
• ie. Ti-47 and Ti-48

47 22
48 22
Ti
Ti
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Atomic Weight

• Average atomic mass for all naturally occurring
  isotopes of an element
• 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
• Shown below element on periodic table

13
Understanding Chemical Reactivity

• Protons define the element
• Neutrons define the isotope of an element
• Electrons control chemical reactivity

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If the protons and neutrons are crammed in the
center of the atom, just where are those
electrons?
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Electrons in an Atom
White Light Source
Hydrogen gas
Beam of light caused by excited electrons in
hydrogen only contained certain wavelengths (or
colors) of light
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Exciting and De-Exciting Electrons
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Bohr Model for Electrons

• Electrons exist in distinct shells about the
  nucleus
• Each shell corresponds to a row in the periodic
  table
• Ignoring the transition metals, shell 1 holds 2
  electrons and each shell after holds 8 electrons
• Shells are filled from the inside out (low energy
  to high energy)
• Each shell out is held a little less tightly by
  the positive nucleus
• Outer shell of electrons controls chemistry
  because easiest to manipulate

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Electron Arrangements
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Reactivity Trends

• Elements in a group react similarly because of
  similar valence (outer most) electron counts

Alkali Metal Trends
Li H2O ? LiOH H2
(Li OH-)
Na H2O ? NaOH H2
(Na OH-)
K H2O ? KOH H2
(K OH-)
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Valence Shell Electrons

• Outermost shell of electrons
• Usually involved in chemical reactions
• Ions are formed by adding or removing electrons
  from outershells
• Typically, all electrons after the last noble gas
• Columns have same valence electrons
• examples C, Si, F, Br
• Play Flame Test Movie

21

• How many valence electrons does oxygen have?
• 2
• 8
• 4
• 6

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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

• 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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Water a Covalent Molecule
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Covalent Bonding Atomic Radius

• Bonding Radii are slightly smaller than
  nonbonding radii
• Bond distance is predicted by sum of bonding
  radii for each atom

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Covalent Bonds

• Two nonmetals come together to make bonds
• Valence Electrons are used to make bonds
• Two atoms share electrons to make a bond
• Each bond (line representation) contains 2
  electrons
• No Formal charge ( or -) on either atom

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Lewis Dot Structures for Elements(Showing
Valence Electrons)
Place valence electrons around an atom symbol,
two on each side
Ne has 8 electrons
Lewis dot Structure
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Bonding Trends (Sometimes Broken)

• Atoms pair up so that each sees at least a full
  valence shell (2 e- for H and 8 e- for other
  atoms)
• To make neutral covalent compounds, atoms make 1
  bond for each electron that it needs to see 8

• C has 4 valence electrons
• C needs 4 electrons to get to 8
• C makes 4 bonds
• N has 5 valence electrons
• N needs 3 electrons to get to 8
• N makes 3 bonds

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Drawing Lewis Structures of Covalent Compounds

• Count total number of valence electrons from all
  atoms in formula
• ie NH3
• Most molecules are symmetric so organize the
  atoms in a symmetric fashion
• If there is only one of an atom, it is probably
  central atom
• Hydrogens go on the outside
• Carbons often make bonds with themselves

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Drawing Lewis Structures of Covalent Compounds

• Place electrons around each atom
• Try to fill outer shells of each atom
• Dont use more electrons than you start with
• Each pair of electrons between atoms makes one
  bond
• If central atoms has not obtained an octet (8
  e-), make multiple bonds by reorganizing
  nonbonding electrons
• ie. NN (double bond)
• ie. N?N (triple bond)
• Check to make sure that you didnt use too many
  or too few electrons

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Formation of Sodium ChloridePlay Movie
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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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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Predicting Ionic Charges
Atoms often gain or lose electrons to form stable
ions that have the same number of electrons as a
noble gases (octet rule, original valence shell
is either full or completely empty)

• 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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Writing Ionic Formulas

• Write cation then anion
• Pair cation and anion up with smallest whole
  number combination to make a neutral salt (look
  for common denominators for overall charge)

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Breaking Ionic Compounds Into Ions

• Look for at least 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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Polyatomic Ions (Covalently Bonded with a Charge)

• Groups of atoms covalently bonded together that
  have an overall charge to help fill the octets
• Memorize common polyatomic ions (names, formulas,
  and charges) p 108
• Whether the polyatomic ion is the cation or the
  anion, use the name give in p 108
• Parentheses are used if more than one polyatomic
  ion is needed

ie. Na3PO4
sodium phosphate
3 Na and 1 PO43-
ie. (NH4)2SO4
2 NH4 and 1 SO42-
ammonium sulfate
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41
Naming Ionic Compounds

• Name cation then anion
• Alkali and Alkaline Earth cations are named with
  the element name
• ie. Na is sodium cation
• ie. Ca2 is calcium cation
• Other Metal cations are named with the element
  name followed by a roman numeral indicating
  charge
• ie. Fe2 is iron(II), Cu is copper(I), and Cu2
  is copper(II)
• Anions are named by taking the root of the
  element name and adding ide
• ie. O2- is oxide anion and Cl- 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. CaS
is calcium sulfide
ie. CuO
is copper(II) oxide