Chemical Foundations

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Chemical Foundations

• Chapter 1

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Scientific Method

• process of studying natural phenomena,
  involving observations, forming laws, and
  theories, and testing of theories by
  experimentation

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Step to the Scientific Method

• Observation use senses
• Qualitative -- descriptive, nonnumeric
• Quantitative -- results in a definite form,
  usually as numbers and units.

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2. Formulating a Hypotheses possible
explanation for what is observed 3. Performing
Experiment running tests to see if hypothesis
is true
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• Theory (Model) a set of tested hypotheses that
  gives a overall explanation of some natural
  phenomenon.
• A theory is an interpretation of why nature
  behaves in a particular way.

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International System of measurements (SI system)

• Based on the metric system and units derived from
  the metric system

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The Fundamental SI Units
Physical Quantity Name of Unit Abbreviation
Mass Kilogram Kg
Length Meter m
Time Second s
Temperature Kelvin K
Amount of Substance mole mol
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Metric Prefixes

• G Giga 109
• M Mega 106
• k kilo 103
• h hecto 102
• da deka 101
• Base 100

• d deci 10-1
• c centi 10-2
• m milli 10-3
• µ micro 10-6
• n nano 10-9
• p pico 10-12

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• SIGNIFICANT FIGURES
• To determine the number of significant figures
  in a written number complete one of the following

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Uncertainty of measurement

• Accuracy How close to the true value you are
• Precision How close your measurements are to
  each other

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• Random Error means a measurement has the
  probability of being high or low this occurs in
  estimating the value of the last digit of a
  measurement.
• Systematic error occurs in the same direction
  each time either always high or always low

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Rules for Counting Sig Figs

1. Nonzero integers
2. Leading zeros ex. 0.0025
3. Captive zeros ex. 1001
4. Trailing zeros ex. 1020 or 1.00
5. Exact numbers infinite number

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• 1. Qualifying statement The numerical value
  has no decimal place written.
• Action to take Count from the first non-zero
  digit to the last non-zero digit.
• Examples
• 18,004 - 5 significant digits
• 10,040,000 4 sig digs
• 10 1 sig dig

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2. Qualifying statement The numerical value
contains a decimal place. Action to take
Count from the first non-zero digit to the end
of the number. Examples 100.00 5 sig
digs 1,050. 4 sig digs 0.000 145 00 5 sig
digs
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Extra Numbers

• Numbers that were NOT obtained using measuring
  devices but were determined by counting
• Ex 10 experiments, 3 apples, 8 molecules

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Practice

1. 10300
2. 1.03 x 1023
3. 0.004050
4. 10 apples

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Rules for math functions

• Multiplication or division
• The answer should have the same number of sig
  figs as the least precise measurement
• Example 1.34 x 0.04 .05

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Rules for math functions

• Addition or Subtraction
• The answer should have the same number of decimal
  places as the least precise measurement
• Example 1.34 2.3 3.6

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Dimensional Analysis

• Converting from one system of units to another
• Sample Convert inches to centimeters (1 in
  2.54 cm)
• 7.O0 inches ________________cm

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Multistep Dimensional Analysis

• Kilometers meters yards miles
• Must have conversion factors for each step.

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Temperature

• K oC 273.15
• or
• oC K 273.15

• oC (oF 32) 5/9
• or
• oF (oC x 9/5) 32

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Practice

• Convert 15oF to co
• Convert 310K to of

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• Density the relationship between an objects
  mass and its volume.
• The denser object will be on the bottom
• Density Mass / Volume (D M / V)
• Example D ? If the mass of an object is 114 g
  and takes up a volume of 10.0 ml.
• D 11.4 g/ml.

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Matter Flowchart
MATTER
no
yes
Can it be physically separated?
Heterogeneous Mixture
Homogeneous Mixture (solution)
Compound
Element
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Pure Substances

• Element
• composed of identical atoms
• EX copper wire, aluminum foil

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Pure Substances

• Compound
• composed of 2 or more elements in a fixed ratio
• properties differ from those of individual
  elements
• EX table salt (Nail)

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Pure Substances

• For example

Two different compounds, each has a definite
composition.
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Mixtures

• Variable combination of 2 or more pure substances.

Heterogeneous
Homogeneous
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Mixtures

• A physical blend of two or more substances.
• Types of Mixtures
• Homogeneous ( Solution)
• Heterogeneous

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Homogeneous Mixture

• One that has a completely uniform composition
• ? components are evenly distributed throughout
  the sample ex. Salt solution
• ? it has a single phase

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Heterogeneous Mixture

• One that is not uniform in composition
• ? has two or more phases
• Ex. Oil and water, salad, milk

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Classification of Matter

• Copper
• Coke
• Pure water
• Sandy water

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States of Matter

• Solid definite shape definite volume
  particles are packed closely together
• Liquid definite volume, but takes the shape of
  the container particles close, but able to move
  around. Liquids flow, almost incompressible and
  expand when heated.

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• Gas indefinite shape volume particles are
  spread far apart. Gases expand to completely
  fill the container

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States of Matter
Property Solid Liquid Gas
Shape Definite Indefinite Indefinite
Volume Definite Definite Indefinite
Expansion on heating Very slight Moderate Great
Compressibility Nil Almost nil Large
Structural arrangement Closely packed Moderate - flows Far apart (move freely)
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Ways to separate a Mixture

1. Distillation separating a solution with
   different boiling pts

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• Filtration separate a solid and liquid mixture
• Chromatography used to separate mixtures by
  their polarity

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• Physical Change
• Chemical Change

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Changes in Matter

• Freezing liquid? solid
• Melting solid ? liquid
• Evaporation liquid ? gas
• Condensation gas ? liquid
• Sublimation solid ? gas

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• Extensive properties physical properties that
  depend on the amount of substance present ex.
  mass, length
• Intensive properties physical properties that
  are not dependent on the amount of substance
  present ex. Density, melting point

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Chapter 1 Questions (p. 32-37)

• 23, 26, 32, 34, 46, 48, 50, 56, 58, 62, 70, 72,
  74, 80, 82, 89

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Atoms, Molecules, and Ions

• Chapter 2

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Law of Conservation of Mass

• Mass is neither created nor destroyed

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Law of Definite Proportion

• Given compound always contains exactly the same
  proportion of elements

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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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Models of Atoms

• Atomic Models
• Chemical properties of atoms, ions, and molecules
  are related to the arrangement of the electrons
  within them.
• John Dalton 1st atomic model considered the
  atom as a solid indivisible mass. Arranged by
  atomic masses or atomic weights.

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Daltons Atomic Theory

• 1. All elements are composed of tiny indivisible
  particles called atoms.
• 2. Atoms of the same element are identical. The
  atoms of any one element are different from those
  of any other element.
• 3. Atoms of different elements can physically
  mix together or can chemically combine.
• 4. Chemical reactions occur when atoms are
  separated, joined, or rearranged. Atoms of
  element are never changed into atoms of another
  element as a result of a chemical reaction.

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• J.J. Thomson revised Daltons model by proposing
  that electrons were stuck to the outside of the
  atom.
• Ernest Rutherford proposed the nuclear atom, in
  which electrons surround a dense nucleus composed
  of protons and neutrons.

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• Chadwick discovered the neutron.
• Niles Bohr student of Rutherford proposed that
  electrons are arranged in concentric circular
  paths (orbits) around the nucleus.

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Atoms tiny particles that make up matter

• Structure of Atoms
• Nucleus center of the atom
• protons positively charged subatomic particles
  that is found in the nucleus dictates the
  identity of the atom

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• Neutron subatomic particle with no charge found
  in the nucleus
• - Discovered by James Chadwick

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• Electron negatively charged found outside the
  nucleus ( electron cloud)
• - Discovered by J.J. Thomson using a cathode ray
  - the rays were attracted to a metal plate of
  positive charge.

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• Atomic Number - the of protons in the nucleus
• - the of protons the of electrons
• Mass Number total of protons neutrons in an
  atom
• - To find the of neutrons subtract the mass
  from the atomic

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• Example Nitrogen (147N)
• Mass number 14
• Atomic number 7
• of protons 7
• of electrons 7
• of neutrons 14 7 7

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• Isotopes - Has the same of protons, but
  different s of neutrons
• Atomic Mass Unit (AMU) 1/12 the mass of
  carbon
• Average atomic mass weighted average of the
  masses of the isotopes of an element

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PERIODIC TABLE

• Periodic Table an arrangement of elements
  according to similarities in their properties
• There are 92 naturally occurring elements.
• Dmitri Mendeleev drew the first periodic table
  Russian chemist arranged the first periodic table
  of elements in 1871. Arranged by atomic mass

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• The periodic table contains chemical symbol,
  atomic number, average atomic mass, physical
  state of each element, group numbers, and
  electron configuration.
• Moseley Later arranged the periodic table by
  atomic number. (Which is the one we use today.)

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MODERN TABLE

• Periods horizontal rows (7 total)
• Groups vertical columns (has similar physical
  chemical properties)
• Metals high electrical conductivity, luster,
  ductile, malleable (Group 1 2A)
• - Alkali Metals Group 1A
• - Alkaline Earth Metals Group 2A

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• Transition Metals Inner Transition Metals
  make up Group B (1B 8B)
• Nonmetals poor conductors, non lustrous
• - Halogens 7A
• - Noble Gases 0
• Metalloids elements that border the stair step
  line
• Group the outermost electrons

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Periodic Trends

• The elements on the periodic table are arranged
  periodically so that trends can be recognized

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Trend of Ions

• 1. You can determine the charge of an ion by
  what group it is in.
• 1A 1 5A -3
• 2A 2 6A -2
• 3A 3 7A -1
• 4A /- 4

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Trend of Electronegativity

• This refers to the ability of an atom to attract
  the electrons of another atom to it.
• Increases across the period (left right)
• Decreases down the group (top bottom)

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Trend of Electron affinity

• Measure of the tendency for atoms to gain
  electrons.
• Increases across the period this is caused by
  the filling of the valence shells
• Decreases down the group this is due to the
  electron entering an orbital far away from the
  nucleus

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Trend of Ionization Energy

• The exact quantity of energy that it takes to
  remove the outermost electron from the atom.
• Factors affecting Ionization Energy
• - nuclear charge
• - distance from the nucleus

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• Ionization energy increases across the period (
  left right) due to increased nuclear charge
• Ionization energy decreases down the group ( top
  bottom)

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Trend of Atomic Radius

• Atomic size is determined by how much space the
  electron takes up. It is also depends on how far
  its valence electrons are from the nucleus.
• The atom will be large if the electron is far
  from the nucleus
• - size increases down a group (top bottom)

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• The atom will be small if the electron is close
  to the nucleus
• - size decreases across the period ( left
  right)
• This is due to an increase in nuclear charge
  pulling them closer the energy level stays the
  same

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Trend of Metallic/Non-Metallic Properties

• Metallic properties elements will form cations
  as they lose electrons (ve charge)
• Non-Metallic properties elements form anions as
  they gain electrons (-ve charge)

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Trend of Melting / Boiling Points

• Melting and Boiling point increase from the right
  side of the periodic table until it reaches
  aluminum and silicon
• Here, melting point and boiling point then begin
  to decrease.

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Trend of Reactivity

• How likely/vigorously an atom is to react with
  other substances
• Metals
• Period decreases from left to right
• Group increases down the group
• The farther left and down you go the easier it is
  for electrons to be taken away. (Higher
  Reactivity)

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Trend of Reactivity

• Non-Metals
• Period increases from left to right
• Group decreases down the group
• The farther right and up you go the higher
  electronegativity vigorous exchange of electrons

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Classification of Elements

• Elements can be classified into 4 groups based on
  electrons.
• 1. Noble gases outermost s p sublevels are
  filled. Belong to group 0. (Also called inert
  gases.)
• 2. Representative elements outermost s or p
  sublevel is partially filled

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• 3. Transition metals metallic elements in which
  the outermost s sublevel and near d sublevel
  contain electrons. (Group B elements)
• 4. Inner transition metals metallic elements in
  which the outermost s sublevel and nearby f
  sublevel generally contain electrons.
  (Lanthanide Actinide series)

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NOMENCLATURE PROCESS USED FOR NAMING CHEMICAL
COMPOUNDS

• Binary Ionic Compounds (criss cross and
  reduce)
• 2. Binary Covalent Compounds (use the prefixes)
• 3. Acids
  (acid rules)

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• Be able to balance and name formulas!

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Chapter 2 Questions (p. 32-37)

• 24, 26, 32, 46, 50, 52, 58, 60, 62, 64, 66, 68,
  70, 72, 74, 78, 87