Introduction to Chemistry Unit 1Chemical Fundamentals

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Introduction to ChemistryUnit 1-Chemical
Fundamentals
Mrs. Dorr
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What is Chemistry?

• Chemistry is the study of matter and its change.
• What is matter? Anything that has mass and takes
  up space.
• mass a measurement of the amount of matter
• Whats the difference between mass and weight???
• Weight measures mass AND Earths gravitational
  pull

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• Examples of matter
• Desk
• You
• Dogs
• Pencil
• Elements
• Examples of non-matter
• love
• Happiness
• Sadness
• Feelings

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How do we describe matter?

• We use physical and chemical properties to
  describe matter.
• Physical property- something that can be observed
  or measured without changing the samples
  composition
• Examples color, density, shape, diameter
• Chemical property- the ability or inability of a
  substance to combine with or change into one or
  more other substances
• Examples reacts with oxygen, does not burn,
  forms rust

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Can these properties be broken down more? Yes!

• Extensive property- dependent on the amount of
  substance present
• Ex- mass, length, volume
• Intensive property- independent of the amount of
  substance present
• Ex- density, melting point, boiling point

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What about the changes of matter, Mrs. Dorr?

• Physical change altering the substance without
  changing its composition
• Chemical change one or more new substances are
  formed (aka- chemical reaction!)
• Indicators of a Chemical Change/Reaction
• 1. formation of a gas (bubbling)
• 2. formation of a solid (precipitate)
• 3. energy is released or absorbed (through
  sound, light, heat)
• 4. color change

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What do we do with all of this? How does
scientific discovery take place?

• Steps of the Scientific Method
• State the Question
• Collect Information (research!)
• Form a Hypothesis
• Experiment (test the hypothesis)
• Observe!
• Record Study Data
• Make a Conclusion

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Measurement Significant Figures

• All measurements consist of a number and a unit
  (with out both the measurement is useless!!!!)
• Types of Units
• Base Unit a unit that is based on an object or
  event in the physical world
• Examples time (seconds), length (meters), mass
  (kilogram)
• Derived unit a unit that is defined by a
  combination of base units
• Examples volume (cm3), density mass / volume

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Density

• Density- the amount of matter per volume
• Note g/mL g/cm3
• Example Determine the density when the mass of a
  metal is found to be 8.1g and it displaces 3.0mL
  of water.
• Example Determine the mass of 0.075L of a
  substance with a density of 3.2g/L.

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Scientific Notation

• Scientific Notation a number between 1 and 10
  raised to a power of 10x
• When multiplying powers of 10? add powers
• When dividing powers of 10? subtract powers

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Examples

• 1. 5.2 x 105
• 2. 5.6 x 10-6
• 3. 0.0027
• 4. 3105.6

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Dimensional Analysis

• a method of problem solving based on the units
  that describe matter
• 1. How many seconds are there in 15.75 hours?
• 2. How many kilometers are there in 50.5 meters

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Metric Conversions Dimensional Analysis

• Convert 750 mL to Liters.
• Convert 3.2 kg to milligrams.
• Convert 0.4 Mm to micrometers.
• Convert 1314 dL to kL.
• Convert 413 cm to hL.

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Accuracy in Measurement

• Accuracy- how close a measurement is to an
  accepted value
• Precision- how close a series of measurements are
  to each other
• error (see side board for eqn.)
• ? Example The accepted value for the density of
  gold is 19.3 g/cm3. A student finds it to be
  19.0 g/cm3. Determine the error

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Significant Figures determine the certainty of
your measurement

• Non-zero integers (2, 353) always count as
  significant figures
• Zeros- 3 classes
• Leading zeros- zeros that preceed all the nonzero
  digits DO NOT COUNT as sig figs. ex- 0.0025
  (how many sig figs?)
• Captive zeros- zeros between nonzero digits
  ALWAYS count as sig figs. ex- 1.0081 (how
  many sig figs?)
• Trailing zeros/ Zeros at the right side of
  numbers are only significant if there is a
  decimal point after it. ex- 5000 vs. 5000.

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• Exact numbers are always significant figures
• Examples on Worksheet

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Operation Rules

• For multiplication division
• The number of sig figs in the result must be the
  same as the least precise measurement in the
  calculation
• Ex. 4.56 x 1.4
• For addition subtraction
• The result must have the same number of decimal
  places as the least precise measurement
• Ex. 12.11 18.0 1.013

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Significant Figure Examples

• 27 / 4.148
• 2.85 3.4621 1.3
• 1.65 x 14
• 7.442 7.429
• 37,100,000 write this w/ 2 sig figs, 3 sig figs
  and 1 sig fig.

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Temperature Conversions

• We will only use the Celsius or Kelvin scale
• The Kelvin scale is based on 0 K (absolute zero)
  and 273 K 0oC
• There are NO NEGATIVE KELVIN VALUES!!!
• To convert to Kelvin from Celsius
• TKelvin TCelsius 273 K
• Example Convert 57 degrees C to Kelvin.
• To convert to Celsius from Kelvin
• TCelsius TKelvin 273K
• Example Convert 400K to degrees C.

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Temperature Conversion Examples

• Convert 35?C to K.
• Convert 155K to Celsius.
• Convert 200?C to K.
• Convert 700K to Celsius.

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Classification of Matter
2 or more pure substances where each substance
retains its chemical properties
a uniform and constant composition
Has mass and takes up space
mixture that is not the same throughout has
different phases
cannot be separated by physical or chemical means
2 or more elements combined chemically
mixture that is the same throughout
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The Periodic Table

• The periodic table is arranged in columns called
  groups and rows called periods
• Elements in the same group have the same number
  of outer/valence electrons
• The elements on the periodic table can be placed
  into 3 categories
• Metals
• Nonmetals
• Metalloids

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Metals

• Metals are on the left side of the table
• Examples Na, Ca, Mg, La, Mn, Fe, Co
• Metals are
• Hard
• Shiny
• Ductile
• Malleable
• Good conductors

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Nonmetals

• Nonmetals are on the right side of the periodic
  table
• Examples C, N, O, Ar, Br, Cl, H, He
• Nonmetals are
• Dull
• Soft
• Brittle
• Poor conductors

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Metalloids

• Metalloids are found on the staircase on the
  right side
• Metalloids include B, Si, Ge, As, Sb, Te, Po, At
• Metalloids are
• Semiconductors
• Possess characteristics of metals and nonmetals
  

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Compounds

• Elements will combine with other elements to form
  a compound
• There are two kinds of compounds
• Ionic compounds
• Molecular compounds
• The key to naming these compounds comes down to
  whether electrons are transferred or shared

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Oxidation Numbers

• The number of electrons in an atom (electrically
  neutral) is equal to the atomic number (protons)
  which is in the upper right corner of an
  elements square.
• All atoms want a full shell of 8 electrons also
  known as noble gas configuration.
• This is called the octet rule

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• For H, He, Li, Be, and B their outer shell is
  full with 2 electrons.
• For ALL OTHER ELEMENTS 8 electrons make up a full
  shell.
• The oxidation number/ charge is equal to the
  number of electrons gained or lost by an atom.
• Metals lose electrons and non-metals gain
  electrons.

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• Atoms that lose electrons are positively charged
  and are called cations.
• Atoms that gain electrons are negatively charged
  and are called anions.
• Write the following elements as ions
• Na
• Ba
• O
• Ga
• P
• Ne
• Si

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Ionic Compounds

• Formed between a metal and a nonmetal
• Ex- NaCl, LiBr, K2O
• Electrons are transferred from the metal to the
  nonmetal
• They are held together by the opposite
  charges/electrostatic attraction between ions

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Rules for naming Ionic compounds

• Name the cation (the positive ion) first and then
  the anion (the negative ion) second.
• If the cation is monatomic (one atom) use the
  full element name
• If the anion is monatomic use the name of the
  element plus the suffix ide.
• If the compound contains a polyatomic ion simply
  use the polyatomic ion name.

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Ionic compounds Examples

• Sodium bonds with fluorine
• Magnesium bonds with chlorine
• Boron bonds with nitrogen

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Writing Formulas

• Identify the oxidation number or charge of each
  element in the compound.
• Write the symbol for each element with the charge
  written in the upper right corner.
• Since the total charge of the compound must be
  ZERO you will use subscripts to make this happen.
• To obtain subscripts? Draw an arrow from the
  charge on the cation/metal down to the bottom
  right of the anion/non-metal symbol. Draw an
  arrow from the charge on the anion/non-metal down
  to the bottom right of the cation/metal symbol.

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Examples- write name formula

• Lithium bonds with fluorine
• Sodium bonds with phosphorous
• Barium bonds with carbon
• Aluminum bonds with chlorine
• Rubidium bonds with silicon

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Subscripts

• Note if the subscripts are the same you dont
  have to write them!
• Ex Beryllium bonds with oxygen.
• BeO

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Try these Write the names and formulas for the
following examples.

• Magnesium bonds with bromine
• Hydrogen bonds with fluorine
• Calcium bonds with oxygen
• Strontium bonds with nitrogen
• Beryllium bonds with iodine

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III. Molecular Compounds

• A molecular compound is formed when two or more
  nonmetals share electrons
• Examples H2O, CH4, CO, CO2, HNO3, CCl4, NH3

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Rules for naming molecular compounds

• The first element in the formula is always named
  first, using the entire element name.
• The second element in the formula is named using
  the root of the element name and adding the
  suffix -ide.
• Use prefixes to indicate the number of atoms of
  each type that are in the compound.

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Prefixes
Mono
hexa
di
hepta
octa
tri
nona
tetra
penta
deca
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Examples

• Name each of the following molecules.
• S4N4
• SeCl2
• H2S
• NO
• N2O
• H2O2
• SiO2

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Writing Formulas from names

• Use the naming rules and prefixes to determine
  the subscripts for each element in the compound.
• Examples Write the formula for the following
  molecular compounds.
• Carbon tetrahydride
• Dihydrogen monoxide
• Diarsenic trioxide
• Dinitrogen tetrachloride

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IV. Polyatomic Ions

• an ion that is made up of one or more nonmetals
  that are covalently bonded together (they share
  e-)

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Ammonium
Acetate
Nitrite
Carbonate
Nitrate
Sulfite
hydroxide
Sulfate
peroxide
cyanide
Permanganate
chromate
Hydrogen carbonate
dichromate
Hydrogen phosphate
Bromate
phosphate
Iodate
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• Because the polyatomic ion exists as a unit,
  NEVER change the subscripts within the ion. If
  you need multiple ions, use parentheses and place
  the subscripts outside of them.
• To name compounds that contain polyatomic ions
  use the ionic compound rules.
• Write the names for these compounds.
• NH4Cl
• HBrO3
• LiCN
• NaOH

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Acids- Two types

• to be an acid it must be dissolved in H2O
• Binary acids made up of hydrogen and one other
  element
• Naming Rules
• Use the prefix hydro- to name the hydrogen part
• Use the root word of the second element name with
  the suffix ic acid
• Examples
• HCl
• HBr
• HF
• H2S

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• Oxyacids made up of hydrogen and an oxyanion (a
  polyatomic ion with oxygen in it)
• Names are made up of the root word of the
  oxyanion and a suffix plus acid.
• If the polyatomic ion ends in ite change the
  suffix to ous acid.
• If the polyatomic ion ends in ate change the
  suffix to ic acid.
• Examples
• H2SO4
• HCN
• HNO2
• HNO3
• HBrO3

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The End!
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