Introduction to Chemistry by 4Bq4d81


									Chemistry Unit Notes
      Science 10 PIB
       Basic Vocabulary

 Matter: Anything that has mass and volume

 Mass: Amount of matter in an object

 Volume: Amount of space an object takes up
              More Vocabulary
 Atoms: Smallest particle of an element that has all the
  properties of matter:
   – Protons- particles in the nucleus with positive charge
   – Electrons- particles orbiting around nucleus with negative charge
   – Neutrons- particles in the nucleus with no charge

 Elements: Simplest form of a pure substance
 Compounds: Two or more elements chemically combined
  to form a new substance
          More Vocabulary

 A mixture is a physical combination of two
  or more kinds of matter.
 A homogeneous mixture is one where the
  components are blended together so well
  that it looks like just one substance e.g.
  salt water or grape juice. Aka a solution
 A heterogeneous mixture is one where
  the different components retain their
  identity and are clearly visible e.g. iron
  filings and salt.
           More Vocabulary
 A pure substance has a definite
  composition which remains the same in
  response to physical changes. E.g.
  copper, water with nothing dissolved in it,
  nitrogen, gold etc.
 Pure substances can be classified into
  elements and compounds. An element
  cannot be chemically separated into
  simpler substances whereas a compound
  can be separated chemically.
Chemical vs. Physical Change
– Physical Change: A change that can occur
  without changing the identity of the
     – Ex. Solid, Liquid, Gas (Phase change)

– Chemical Change: Process by which a
  substance becomes a new and different
     – Ex. Fire
           Sub-Atomic Particles

Part of       Charge       Location      Mass/Size

Electron      - negative   outside       .0006 amu
                           nucleus       (too little to count)

Proton        + positive   inside nucleus 1 amu

Neutron       no charge    inside nucleus 1 amu
Periodic Table
     Using the Periodic Table
             Atomic Number
              – Equal to # protons = # electrons
Cl            – Periodic Table is arranged by this number


             Symbol
              – “Shorthand” for the element – Note 2nd letter is
                always lowercase

             Atomic Mass Number
              – Total AVERAGE mass of Protons + Neutrons +
                Electron Energy Levels
 Electrons are arranged in “Shells” around nucleus in
  predictable locations
 Fill “seats” closest to nucleus first (concert – best seats)
 “Seats” available
    –   Shell    #1                   2 electrons
    –   Shell    #2                   8 electrons
    –   Shell    #3                   8 electrons
    –   Shell    #4                   18 electrons
    –   Shell    #5                   18 electrons
    –   Shell    #6                   32 electrons

   Ex. Carbon has 6 total electrons so…

                          Two electrons on first energy level

                                       Four electrons on second energy level
        Question: Could we fit more electrons on the second energy level if there were more electrons in carbon??
                  Atomic Structure
                         Total # of protons and electrons (in a neutral atom)
    17                            17 protons in nucleus
                                  17 electrons orbiting nucleus
   35.5                 Element Name

          Total Mass of Nucleus
                   36 - 17 = 18 neutrons

(Round Atomic Mass)
                                      Notice: electrons follow energy level rules
                                      from previous slide.
      Atomic Mass – Fractions?
 Look at Chlorine (atomic number 17)
 Atomic mass of 35.5? I dont’ get it!
 Where does the 35.5 come from?
    – 0.5 protons? 0.5 neutrons?  No

   Atomic Mass = average number of protons
    and neutrons in nature
              More Practice
   Determine the name, number of protons,
    neutrons and electrons for each element
    shown and draw…

                      8              26

     P               O              Fe
                     16             56

   An isotope is a variation of an element
    (same protons) but can have diff. # of

   Ex: carbon (atomic mass = 12.011)
    – Carbon (14) and carbon (12) exist in
 Change in electrons which gives an atom a
  charge (+ or -)
 You can only add or subtract electrons!
      (protons don’t change)
    – Ex.   Count the number of electrons below…

                       Carbon ion (-1 charge)      Neutral Carbon    Carbon ion (+1 charge)
                       7 electrons (-)             6 electrons (-)   5 electrons (-)
                       6 protons (+)               6 protons (+)     6 protons (+)
              Valence Electrons
   An electron on the outermost energy shell of an
   Important to understand because this is a key factor
    in how atoms will BOND with each other
   Octet rule – stable atom will have 8 electrons in that
    outer shell
   Practice – Valence # of
     – Chlorine?
     – Neon?
     – Nitrogen?
     – Oxygen?
         Electron Dot Diagrams
 a diagram that represents the # of valence
  electrons in an atom of an element.
 The amount of electrons is displayed by dots
  around the symbol of the element.
 Ex.

           Types of Chemical Bonds
   Ionic-     Two elements bond by transferring electrons to create ions
    that attract together (+ is attracted to - after an electron is transferred)

   Covalent-        Two elements bond by sharing electrons (strongest
    bond type)

   Metallic-       Two metals bond and form a “common electron cloud”.
    This is a cluster of shared electrons (weakest bond type)
   Examples of Bonding
                             Predicting Bonds
   Ionic Bond = metal to non-metal
   Covalent = non-metal to non-metal
   Metallic = metal to metal

Do you understand why? HINT: the numbers at the top of the table indicate the # of valence electrons for each column
                 Oxidation Numbers
   Oxidation numbers are assigned to each element

   They represent a predicted “charge” of an atom/ion
    when it bonds with another element.
         (tells us if the atom would prefer give or take electrons, and how many).

   They help us to predict what compounds will form
    when two elements get together.

   Oxidation numbers are labeled like this:
         Na 1+
         O 2-
    How to Use Oxidation Numbers
Oxidation Number indicates the number of electrons lost, gained or shared when
   bonding with other atoms.

    Ex. Na wants to lose an electron. If an electron is lost, it
      becomes a +1 charge

    SO: oxidation number for Na = 1+

    Ex. Cl wants to gain an electron. If an electron is gained, it
      becomes a -1 charge
       SO: oxidation number for Cl = 1-
           Oxidation Numbers

   Each column going down the periodic table
    has elements with the same oxidation
   Label the oxidation numbers on your periodic table at the top of each
    column as shown here:

     1+ 2+                                            3+ 4(+/-) 3- 2- 1- 0
       Rules for using oxidation
     numbers to create compounds
1. Positive ions can only bond with negative ions and vice
2. The sum of the oxidation numbers of the atoms in a
    compound must be zero (the key is to stay balanced)

3. If the oxidation numbers are not equal to zero, then you
    must add additional elements until they balance at zero.

4. When writing a formula the symbol of the Positive (+)
   element is followed by the symbol of the negative (-)
  Examples of Forming Compounds
Ex. Na (+1) + Cl (-1) = NaCl
   Are these oxidation numbers already equal to zero?
   If so, you don’t need to add any extra elements to combine them into a compound, so the answer is
   simply NaCl

Ex. H (+1) + O (-2) = H2O
   How many +1 would you need to balance the -2 to zero?
   Since you need 2 atoms of the 1+ to balance the 2- to zero the resulting compound would be H2O
   In other words: to combine H with O, you MUST have 2 H to balance the oxidation numbers to zero
   2+ and 2- = ZERO

Ex. Al (+3) + S (-2) = Al2S3
   This one is tricky…we are not even close to balancing + and - to zero.
   Because of this we must have more than one Al and more than one S in our final equation.
   By using 2 Aluminums instead of just1 we would have 6+
   By using 3 sulfers instead of just 1 we would have 6-
   Since these are now equal to zero, we combine 2 Aluminums and 3 Sulfers to make Al 2S3
           Chemical Reactions
   Chemical Reaction: a process in which the
    physical and chemical properties of the
    original substance change as new
    substances with different physical and
    chemical properties are formed
        Chemical Reaction Basics
                H2 + O2 --> H2O
                  Reactants          Products

Reactants- substance that enters into a reaction

Products- substance that is produced by a chemical reaction
    Evidence of Chemical Change
        EPOCH is an acronym that stands for evidence that a
         chemical reaction has occurred.

E–    Effervescence (bubbles and/or gives off gas)
P –   Precipitate (solid crystals form)
O–    Odor (change of smell is detected)
C –   Color change
H –   Heat (reaction either heats up or cools down)

        Does sighting evidence of a chemical reaction mean that a
         chemical reaction has undoubtedly taken place?
         Types of Reactions
            Romance Chemistry :)

Synthesis- Marriage/Dating
     A + B = AB
Decomposition- Divorce/Breakup
     AB= A + B
Single-Replacement- Dance Cut In
     A + BC = AC + B
Double-Replacement- Dancing couples
  switch partners.
     AB + CD = AC + BD
Cartoon Chemistry

  This is an example of synthesis
Cartoon Chemistry

This is an example of a decomposition
Cartoon Chemistry

 This is an example of a single replacement
Cartoon Chemistry

This is an example of a double replacement
       Reaction Types Review…
   Match each chemical reaction with one of
    the reaction types on your chemical
    – Zn + 2HCl  H2 + ZnCl2
    – N2 + 3H2  2NH3
    – 2KI + Pb(NO3)2  2KNO3 + PbI2
    – 2MgCl  Mg2 + Cl2
       Conservation of Mass
 Atoms cannot be created or destroyed in a
  chemical reaction.
 What goes in must come out.
 So we must balance equations to conserve
              Balancing Equations
   Rules:
    – We can not add or subtract subscripts from either
      side of the equation
    – We can only add coefficients to the front of each
   Ex.      2H2 + O2 --> 2H2O
                H=4                     H=4
                O=2                     O=2

                  Before   must match    After

             See “Balancing Act” worksheet for more examples…
       Solution Chemistry
 Mixtures: Matter that consists of two or more substances mixed
  but not chemically combined
 Solutions: Homogeneous Mixture in which one substance is
  dissolved into another
        Solute = Substance that gets dissolved (ex. Kool-Aid powder)
        Solvent = Substance that does the dissolving (ex. Water)
 Acid: Compound with a pH below 7 that tastes sour and is a
  proton donor.
        Ex. Citrus foods
 Base: Compound with a pH above 7 that tastes bitter and is a
  proton acceptor
        Ex. Cleaning Products (soap)
                 Acids and Bases
-   Solutions can be acidic or basic

-   Acids and Bases have unique properties when dissolved in
     - Acids = sour taste
     - Bases = bitter taste

-   Indicators are substances that change color when mixed
    with a solution, which helps to determine if a substance is an
    acid or a base. (pH paper, Litmus paper, cabbage juice)
 Proton donors (H+)
 Acids contain hydrogen and produce positive
  ions (H+) when dissolved in water
 Acids = good electrolytes
 Examples of acids:
    –   Lemon Juice
    –   Citric Acid
    –   Carbonic Acid
    –   HCl
 Proton acceptors
 Bases contain hydroxide ions (OH-) when
  mixed with water.
 Bases = weak electrolytes
 Examples of bases:
    – Ammonia
    – Soap
    – Bleach (chlorine)
   Combining Acids and Bases
-Mixing acids and bases is a balancing act.
 (like a teeter-totter)

Acid + Base = neutral (water and salt)
   Combining Acids and Bases
Acid + Base = neutral (water and salt)

  H+    + OH-  HOH + Salt
 Acid     Base  water

Ex. HCl +     NaOH        H2O + NaCl
    Measuring Acids and Bases
 pH scale- used to measure the acidity of a
 Measure pH with indicators
 pH scale goes from 0 – 14
 0 = very acidic
 14 = very basic
 7 = neutral
Acids and Bases

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