3.1 Early Theories of the Atom

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Chapter 3
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3.1Early Theories of the Atom

• Four men paved the road to how we see the atom
  today
• Democritus
• Aristotle
• John Dalton
• Rutherford (Chapter 3.2)
• Xenon on Copper Atoms

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Democritus

• Matter is composed of empty
• Space through which atoms move.
• Atoms are solid, homogeneous, indestructible, and
  invisible.
• Different kinds of atoms have different sizes and
  shapes.
• The differing properties of matter are due to the
  size, shape, and movement of atoms.
• Apparent changes in matter result from changes in
  the groupings of atoms and not from changes
  themselves.

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Aristotle

• One of the most influential
• philosophers.
• Wrote extensively on many subjects, including
  politics, ethics, nature, physics, and astronomy.
• Most of his writings have been lost through the
  ages.

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John Dalton

• All matter is composed of extremely small
• particles called atoms.
• All atoms of a given element are identical,
  having the same size, mass, and chemical
  properties. Atoms of a specific element are
  different from those of any other element.
• Atoms cannot be created, divided into smaller
  particles or destroyed.
• Different atoms combine in simple whole-number
  ratios to form compounds.
• In a chemical reaction, atoms are separated,
  combined, and rearranged.

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Yesterdays Scientists in Today's Labs

• Daltons Atomic Theory was a huge step toward
  defining the atom, and learning about it.
• Today scientists have used his theory as
  machinery and microscopes become better to find
  wrong statements in his theory.
• Sodium, on Iodine on Copper atoms

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Flaws in Daltons and Greek Theory

• Atoms are not invisible, just invisible to the
  naked eye. With a powerful microscope, scientists
  can rearrange atoms into different shapes.
• Atoms of an element have a slightly different
  mass.
• Atoms can be split into sub-particles

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Defining the Atom

• After Daltons watershed event in Chemistry,
  scientists wanted to study the concept of the
  atom. In the next section, you will learn the
  structure of the atom, and what it is made of.

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3.2Subatomic Particles and the Nuclear Atom

• Rutherfords Gold Foil Experiment
• Subatomic Particles
• -Electrons
• -Neutrons
• -Protons

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Rutherfords Gold Foil Experiment

• The Plum Pudding Model stated that negatively
  charged electrons were distributed throughout a
  uniform positive charge.

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Gold Foil

• Ernest Rutherford began an experiment named the
  Gold Foil Experiment, where he set up a narrow
  beam of alpha particles aiming at a coated screen
  surrounding the gold foil. A zinc sulfide coated
  screen provided a flash of light whenever an
  alpha particle struck it.

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Gold Foil (What was expected)

• Rutherford expected most of the fast moving and
  relatively massive alpha particles to pass
  straight through the gold atoms.

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Gold Foil (What Happened)

• Instead of the alpha particles going through,
  some shot back, others were altered off of course.

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Gold Foil (The Reason)

• Rutherfords nuclear model of the atom explains
  the results of the gold foil experiment. Most
  alpha particles pass straight through, being only
  slightly deflected by the electrons, if at all.
  The strong force of repulsion between the
  positive nucleus and the positive alpha particles
  causes the large deflections.

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• Its as if you shot a cannon ball at a piece of
  tissue paper and it came right back at you!
• -Rutherford

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Electrons

• Negatively charged particles that are part of all
  forms of matter.
• Located in the space surrounding the nucleus.
• Relative electron charge of 1-
• Relative mass 1/1840
• Actual mass in grams 9.11 x 10-28

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Protons

• A proton is a subatomic particle carrying a
  charge equal to but opposite that of an electron.
  
• It has a positive charge.
• Located in the nucleus.
• Relative electron charge of 1
• Relative mass 1
• Actual mass 1.673 x 10-24

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Neutrons

• A neutron has a mass equal to that of a proton,
  but it carries no electrical charge.
• It is located in the nucleus.
• Relative Electrical charge of 0
• Relative mass 1
• Actual mass 1.675 x 10-24

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4.3How Atoms Differ

• Atomic Number
• Isotopes
• Mass Number
• Atomic Mass
• The Periodic Table of Elements

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Atomic Number

• Atomic Number is the number of protons in an
  atom.
• Atomic Numbernumber of protonsNumber of
  electrons
• Helium has an atomic number of two, therefore it
  has 2 electrons and 2 protons

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Isotopes

• Elements with the same number of protons, but a
  different number of neutrons are called isotopes.
• They differ in mass, and those containing more
  neutrons have a bigger mass.

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Mass Number

• The mass number represents the sum of the number
  of protons and neutrons in an elements nucleus.
• Potassium-39 isotope has 19 protons and 20
  neutrons, therefore its mass number is 39
  (192039)
• Number of neutrons mass number atomic number

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Atomic Mass

• The atomic mass of an element is the weighted
  average mass of the isotopes of that element.
• Ex) Chlorine has a mixture of 75 chlorine-35,
  and 25 of chlorine-37. The actual atomic mass is
  found by the sum of the products of each isotopes
  percent abundance times its atomic mass.

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Platinum Iron on Copper Nickel Carbon
Monoxide on Platinum
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Chapter 5
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Chapter 5

• Light and Quantized Energy
• Classification of Elements
• Electron Configurations

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Chapter 5 (Insights)

• Wavelength
• Frequency
• Amplitude
• Electromagnetic Radiation
• Electromagnetic Spectrum
• Other Vocab

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Wavelength

• Represented by , is the shortest
  distance between equivalent points on a
  continuous wave. Wavelength is measured from
  crest to crest or from trough to trough.
  Wavelength is usually expressed in meters,
  centimeters, or nanometers
• (1nm1 X 10-9)

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Wavelength
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Frequency

• Represented by v , is the number of waves
  that pass a given point per second. One (Hz), the
  SI unit of frequency, equals one wave per second.
  The Hz stands for hertz.

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Amplitude

• Is the waves height from the origin to a crest,
  or from the origin to a trough.

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Electromagnetic Radiation

• A from of energy that exhibits wavelike behavior
  as it travels through space.
• Ex) gamma rays
• alpha rays
• beta waves

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Electromagnetic Spectrum

• The EM spectrum, encompasses all forms of
  electromagnetic radiation, with the only
  differences in the types of radiation being their
  frequencies and wavelengths.

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Other Vocab

• Quantum-The minimum amount of energy that can be
  gained or lost by an atom.
• Photon-A particle of electromagnetic radiation
  with no mass that carries a quantum of energy.
• Atomic Emission Spectrum-the finger print of an
  atom caused by the light it emits.
• Heinsenburg Uncertainty Principle-The exact
  momentum and position of an e- can never be
  precisely known.

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Quantum Theory and the Atom

• Energy sublevels-The atomic orbitals represent
  the electrons probability clouds of an atoms
  electron. All s-orbitals are spherical in shape
  and increase in size with increasing principle
  quantum number. The p-orbitals are dumbbell
  shaped, and each are related to an energy
  sublevel that has equal energy. They can contain
  up to 6 electrons.

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Principles

• Aufbau Principle-Each electron occupies the
  lowest energy orbital available.
• The Pauli exclusion Principle-The maximum of two
  electrons may occupy a single atomic orbital.
• Hunds Rule- The single electrons with the same
  spin must occupy each equal energy orbital before
  additional electrons with opposite spins occupy
  the same orbitals.

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Chapter 6

• Development of the Modern Periodic Table
• Classification of Elements
• Periodic Trends

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John Newlands

• Organized elements by mass and noticed a
  repeating trend every eighth element. He came up
  with the Law of Octaves that was frowned upon
  by other scientists, because music and science do
  not mix.

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Newlands' Octave Arrangement
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Dmitri Mendeleev

• Russian scientist that organized the periodic
  table between atomic mass and properties and was
  able to predict unfound elements of his time. His
  table was widely accepted.

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Orbital Structureand Electron Configurations
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The S Orbital

• Every element has a 1sx, the x number depends on
  how many electrons are there. Hydrogen has one
  electron, therefore, I is 1s1. Helium has 2
  electrons, therefore it is 1s2. Then the s
  orbital is filled, so to go onto Lithium, another
  s orbital is needed, the 2s orbital, when
  Beryllium fills that one, a p orbital takes the
  next electron to be Boron 1s2, 2s2, 2p1
• (Remember, a p orbital can hold 6 electrons)
• Neon10 1s2, 2s2, 2p6

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The P Orbital

• A p orbital hold six electrons, when a p shell is
  filled, ex) Neon, the outer most electrons in the
  outer most shell are called valence electrons.
  When Neon fills the 2p orbital, the 4s orbital
  begins again. Na, Sodium is labeled as Ne 3s1
  (This is called a Noble Gas notation).

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Groups of the Periodic Table

• Chapter 7
• Review notes from class on periodic groups and
  the class presentations.
• Halogens, Alkaline earth metals, alkali metals,
  Noble gasses, d-block, and f-block are key groups.

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Chapter 8

• Forming Chemical Bonds

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8.1 Forming Chemical Bonds

• Chemical Bond-The force that holds two atoms
  together may form by the attraction of a
  positive nucleus for negative electrons.
• Cation- An ion that has a positive charge forms
  when valance electrons are removed, giving the
  ion a stable electron configuration.
• Anion- An ion that has a negative charge forms
  when electrons are added to the outer shell,
  giving the ion a stable electron configuration.

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8.2 The Formation and Nature of Ionic Bonds

• Electrolyte- An ionic compound whose aqueous
  solution conducts an electron current.
• Lattice Energy- The energy required to separate
  one mole of the ions of an ionic compound, which
  is directly related to the size of the ions
  bonded and is also affected by the change of ions.

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8.3 Names and Formulas for Ionic Compounds

• Formula Unit- The simplest ratio of ions
  represented in an ionic compound.
• Monatomic ion- An ion formed from one atom.
• Oxidation Number- The positive or negative charge
  of a monatomic ion.
• Polyatomic ion- A ion made up of two or more
  atoms bonded together that acts as a single unit
  with a net charge.
• Oxyanion-A polyatomic ion composed of an element,
  usually a non metal, bonded to one or more oxygen
  atoms.

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8.4 Metallic Bonds and Properties of Metals

• Electron Sea Model-Proposes that all metal atoms
  in a metallic solid contribute their valance
  electrons to form a sea of electrons, and can
  explain properties of metallic solids, such as
  malleability, conduction, and ductility.
• Delocalized Electrons- The electrons involved in
  metallic bonding that are free to move easily
  from one atom to the next throughout the metal
  and are not attached to a particular atom.

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8.4 Continued

• Metallic bond- The attraction of a metallic
  cation for delocalized electrons.
• Alloy- A mixture of elements that has metallic
  properties most commonly forms when the elements
  are either similar in size (Substitutional alloy)
  or the atoms of one element are much smaller that
  the atoms of the other (interstitial alloy).

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Chapter 9

• Covalent Bonding

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9.1 The Covalent Bond

• Covalent Bond- A chemical bond that results from
  the sharing of valance electrons.
• Molecule- Forms when two or more atoms covalently
  bond and is lower in potential energy than its
  constituent atoms.
• Lewis Structure- A model that uses electron-dot
  structures to show how electrons are arranged in
  molecules. Pairs of dots or lines represent
  bonding pairs.

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9.1 continued

• Sigma Bonds- A single covalent bond that is
  formed when an electron pair is shared by the
  direct overlap of bonding orbitals.
• Pi Bond- A bond that is formed when parallel
  orbitals overlap to share electrons.
• Endothermic- A chemical reaction in which a
  greater amount of energy is required to break the
  existing bonds in the reactants than is released
  when the new bonds form in the product molecules.
• Exothermic- A chemical reaction in which more
  energy is released than is required to break
  bonds in the initial reaction.

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• Sigma Bond
• Pi Bond

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9.2 Naming Molecules

• Oxyacid- Any acid that contains hydrogen and an
  oxyanion.

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9.3 Molecular Structures

• Structural Formula- A molecular model that uses
  symbols and bonds to show relative positions of
  atoms can be predicted for many molecules by
  drawing the Lewis Structures.
• Resonance- Condition that occurs when more than
  one valid Lewis Structure exists for the same
  molecule.
• Coordinate Covalent Bond- Forms when one atom
  donates a pair of electrons to be shared with an
  atom or ion that needs two electrons to become
  stable.

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9.4 Molecular Shape

• VSEPR Model- Valance Shell Electron Pair
  Repulsion Model which is based on an arrangement
  that minimizes the repulsion of shared and
  unshared electrons around the central atom.
• Hybridization- The process by which the valance
  electrons of an atom are rearranged to form four
  new, identical hybrid orbitals.

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9.5 Electronegativity and Polarity

• Polar Covalent- A type of bond that forms when
  electrons are not shared equally.

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Chapter 19

• Acids and Bases

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19.1

• Acidic Solution- Contains more hydrogen atoms
  than hydroxide ions.
• Basic Solution- Contains mor hydroxide ions than
  hydrogen ions.
• Arrenius Model- A model of acids and bases
  states than an acid is a substance than contains
  hydrogen and ionizes to produce hydrogen ions
  aqueous solution and a base is a substance that
  contains a hydroxide group and dissociates to
  produce a hydroxide ion in aqueous solution.

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19.1 Continued

• Bronsted-Lowry Model- A model of acids and bases
  in which an acid is a hydrogen-ion donor and a
  base is a hydrogen-ion acceptor.
• Conjugate Acid- The species produced when a base
  accepts a hydrogen atom from an acid
• Conjugate Base- The species produced when an acid
  donates a hydrogen ion to a base.

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19.1 Continued

• Conjugate acid-base pair- Consists of two
  substances related to each other by the donating
  and accepting of a single hydrogen atom.
• Amphoteric- Describes water and other substances
  that acts as both acids and bases.

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19.2

• Strong acid- An acid that ionizes completely in
  aqueous solution.
• Weak acid- An acid that ionizes only partially in
  dilute aqueous solution.
• Acid ionization Constant- The value of the
  equilibrium constant expression for the
  ionization of a weak acid.

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19.2 Continued

• Strong base- A base that dissociates entirely
  into metal ions and hydroxide ions in aqueous
  solution.
• Weak base- A base that ionizes only partially in
  dilute aqueous solution to form the conjugate
  acid of the base and hydroxide ion.
• Base ionization constant- The value of the
  equilibrium constant expression for the
  ionization of a base.

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19.3

• PH- The negative logarithm of the hydrogen ion
  concentration of a solution acidic solutions
  have PH values between 0 and 7, basic solutions
  have values between 7 and 14, and a PH of 7 is
  neutral (Water).
• POH- The negative logarithm of the hydroxide ion
  concentration of a a solution a solution with a
  POH above 7.0 is acidic and below 7.0 is basic. A
  solution with a POH of 7.0 is neutral.

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Chapter 10

• Chemical Reactions

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10.1 Reactions and Equations

• Chemical Reaction- The process by which the atoms
  of one or more substances are rearranged to form
  different substances. Indicated by changes in
  temp, color, odor, or physical state.
• Reactant- The starting substance in a chemical
  reaction.
• Product- A substance formed in a chemical
  reaction.

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10.1 Continued

• Chemical Equation- A statement using chemical
  formulas to describe the identities and relative
  amounts of the reactants and products involved in
  a chemical reaction.
• Coefficient- In a chemical equation, the number
  written in front of a reactant or product (used
  to balance the equation).

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10.2 Classifying Chemical Reactions

• Synthesis Reaction- A chemical reaction in which
  two or more substances react to yield a single
  product.
• ( A B AB)
• Combustion Reaction- A chemical Reaction that
  occurs when a substance reacts with oxygen,
  releasing energy in the form of light and heat.
• CH4 O2 ---gt CO2 H2O

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10.2 Continued

• Decomposition Reaction- A chemical reaction that
  occurs when a single compound breaks down into
  two or more elements or new compounds.
• AB ---gt A B
• Single-reactant Reaction- A chemical reaction
  that occurs when the atoms of one element replace
  the atoms of another element in a compound.
• AX Y ---gt YX A
• Precipitate- A solid produced during a chemical
  reaction in a solution.