Bonding and Chemical Formulas

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Bonding and Chemical Formulas

• Unit IVA - .2 Chemistry
• October 2006- Revised Nov. 2007

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Bonding

• There are two major types of bonding
• Ionic and covalent
• Covalent bonds occur when two atoms SHARE valence
  electrons
• Ionic bonds are due to a very strong
  electrostatic attraction between two ions (atoms
  who just EXCHANGED electrons)
• In other words, one atom stole electrons from the
  other

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Ions

• Lets start by deciding what an ion is
• An ion is an atom which has either gained or lost
  an electron
• When an atom gains an electron, the ion is
  negatively charged
• Why?
• When an atom loses an electron, the ion is
  positively charged
• Why?

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Ions

• A positively charged ion is referred to as a
  cation
• A negatively charged ion is called an anion
• Metals tend to be cations because they will
  frequently lose electrons
• Nonmetals tend to be anions because they will
  take on more electrons

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Electronegativity

• The difference in electronegativities between two
  atoms, determines whether an ionic or covalent
  bond will form.
• So, what is electronegativity?
• Electronegativity is defined as
• an atoms pull on its electrons
• As a rule of thumb, electronegativity increases
  as you go and to the

Right -gt
UP
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Electronegativity

• Look at your table of electronegativities
• Metals tend to have very LOW electronegativity
  values and NonMetals tend to have quite HIGH
  values
• The difference between any two atoms will decide
  whether an ionic or a covalent bond will form
• If the diff is gt2, then an ionic bond will form
• If the diff is lt1.69, then a covalent bond will
  form
• (between 1.7-2, if M-NM then ionic and if 2 NM
  then covalent)

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Ionic vs Covalent

• What type of bond will form between each of the
  following?H and Cl?
• C and O?
• Na and Br?
• Be and F?
• What trend do you see?
• Typically nonmetals form ionic bonds with metals
  and they form covalent bonds with other nonmetals

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Valence Electrons

• As you recall, the electrons in the outermost
  shell or energy level are called valence
  electrons
• These valence electrons are typically the only
  ones involved with bonding
• Lets use a technique called electron dot
  structures to represent the valence electrons of
  an atom (also called Lewis Dot structures)

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Valence Electrons

• Each dot represents an electron
• First, determine how many valence electrons an
  atom has
• Where will you find this information?
• Then write the chemical symbol for that atom
• Place the correct number of dots around each of
  the 4 sides of the atomic symbol
• Be sure to only put one dot per side until all
  sides are 1/2 filled, then you can start
  doubling up
• Why?

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Electron Dot Structures

• Write the electron dot structure for each of the
  following

• Na
• C
• O
• Cl
• Kr
• Mg

• Li
• B
• P
• S
• Ne
• Ca

• Notice the patterns that will occur within
  families

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Electron Configurations for Ions

• We already know how to write the electron
  configuration for an atom, lets apply it to ions
• We do it the exact same way except we take into
  account the electrons that have been added or the
  missing electrons
• The most stable form of the ion will be that
  which shares the electron configuration with a
  noble gas

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Electron Configurations for Ions

• Noble gases all have the same ending to their
  electron configuration, s2 p6, giving them 8
  electrons in their outermost energy level
• This is where the octet rule comes from
• We know that noble gases are the most stable
  elements, scientists gathered that the reason for
  this is because they have 8 electrons in their
  valence shell

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Electron Configurations for Ions

• What is the noble gas configuration for the
  following ions?
• Ca2
• O 2-
• Li1
• Al3
• H-1

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Electron Configurations for Ions

• What is the noble gas configuration for the
  following ions?
• Ca2 Ar
• O 2- Ne
• Li1 He
• Al3 Ne
• H-1 He

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Ionic Bonding

• Ionic bonding involves the TRANSFER of electrons
  or stealing of electrons as we mentioned before
• Do you recall how we determine which element is
  going to become a cation and which will become an
  anion?

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Ionic Bonding

• Each of you will be given a card with an
  elemental symbol on it
• For that symbol, determine the number of valence
  electrons (this is how many dots you would put
  around the symbol in your Lewis dot structure)
• Then, come up and grab 1 bill for each valence
  electron you have

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Ionic Bonding

• Using the list of compounds, find your partner
  and EXCHANGE electrons accordingly
• Once you have exchanged electrons, you are
  transformed into your respective ions
  (cation/anion)
• You are electrostatically attracted to each other
  so BOND (stand tightly shoulder to shoulder
  until I come around and check your compound)

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Properties of Ionic Compounds

• A solid ionic compound is called a salt
• All salts share 5 characteristics
• 1. Made of crystals
• 2. Conduct electricity
• 3. Have high melting and boiling points
• 4. Are hard
• 5. Are brittle

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1. Made of Crystals

• Attraction between opposite charged ions is SO
  great that there is more than one bond
• A tightly packed cluster of repeating units forms
  the crystal structure
• Ex. NaCl is the formula
• unit for table salt
• crystals

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2. Conducts Electricity

• Electricity needs charged particles that are free
  to move (in solution)
• In salt water, the particles spread out and can
  carry electricity from ion to ion throughout the
  solution
• Electrolytes - ions in solution which carry an
  electric current

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3. High BP and MP

• Because of the strong attractions between the
  oppositely charged ions, it takes a lot of energy
  to break up the particles
• BP - boiling point is the temperature at which
  you have a phase change from a liquid to a gas
• MP - melting point is the temperature at which
  you have a phase change from a solid to a liquid

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4. Hard and 5. Brittle

• Salts are hard due to the strong attraction of
  opposite charges and the layering of crystals
• Salts are brittle, or break up to make a powder
• Layers usually line up so that - and alternate,
  but added energy or pressure can cause to be
  next to and - to be next to -.
• Does it like itNO so it breaks apart into
  powder

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Hydrates

• Some salts can hold water molecules between their
  bonds
• These are called Hydrated Salts
• Possible Uses drying agents or moisture
  indicators
• Ex. CuSO4 . 5 H2O
• In this hydrate, for every salt unit of Copper
  sulfate, 5 molecules of water are trapped

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Percent Composition of Hydrates

• What of CuSO4 . 5 H2O is water?
• First lets find the formula mass of copper
  sulfate by itself
• 63.5 32.1 64.0 159.6 g/mol
• What about the water that is trapped?
• 1.0 (10) 16.0 (5) 90.0 g/mol
• Total molar mass 249.6 g/mol
• mass of water 90.0 / 249.6 x 100 36

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Try another hydrate problem

• Calculate the percent water in NiCl2. 6H2O ?
• NiCl2 129.7 g/mol
• 6 H2O 108.0 g/mol
• Total mass 237.7 g/mol
• 108.0 / 237.7 x 100 45

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Percent Composition

• The law of definite proportions refers to the
  chemical make-up of ONE compound
• Within that compound, the proportion or ratio of
  one element to another will remain the same no
  matter how much of the compound is present
• The law of multiple proportions compares the
  compositions of two different compounds which
  contain the same elements

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Percent Composition

• The Law of definite proportions explains why we
  can have a formula unit for a compound
• This is the simplified version of the elemental
  ratios within the compound
• For example, when joining Mg and Cl what is the
  smallest whole number ratio that can be used to
  join these two together
• Make sure that the charges balance out

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Percent Composition - (refresher)

• If I have an ionic compound of MgCl2, what
  percentage of the whole mass does Mg make up?
• First, find the atomic mass of Mg
• Then, find the atomic mass of Cl and double it
  because there are two atoms of Cl
• What is the total mass?
• Divide the mass of the Mg by the total mass and
  multiply by 100. This is the percent of the
  whole that Mg makes up.

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Percent Composition

• Calculate the percent composition of water in the
  following hydrate
• (NH4)2SO4 . 5 H2O

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Mole Ratios and Hydrate Predictions

• For our hydrated salt lab, you took the mass of
  the hydrated salt before heating
• You then heated it up for 10-15 minutes as
  instructed and took the mass of the anhydrous
  salt
• Anhydrous salt salt without the water
• To determine the amount of water that was in your
  original hydrated salt sample, subtract the mass
  of anhydrous salt from the mass of hydrated salt

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• Now you can do two things
• 1. determine the percent of the salt sample that
  was water
• Do this by taking the mass of water and dividing
  by the total mass of the hydrated salt (before
  heating)
• Multiply by 100 and this is your experimental
  water.

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• 2. The second thing you can calculate is the
  of waters trapped in the salt per formula mass
  unit.
• Do this by first converting your number of grams
  of anhydrous salt to moles (use the molar mass of
  the salt)
• Then convert your mass of water to moles (using
  the molar mass of water)
• Now divide both numbers (of moles) by whichever
  is smaller to get a ratio.
• Clue the ratio will by 1 to ___ (rounded to
  the nearest whole number).
• If the ratio is very close to a HALF number, then
  double both. Ex. 1 2.5 gets doubled to 25
  ratio

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Hydrate Prediction Example

• In the lab, you measure out 5.25g of barium
  chloride (BaCl2 )
• After heating, the mass is 4.50g.
• Calculate the amound of water that was in the
  hydrated salt.
• 5.25g 4.50g
• 0.75 g
• Calculate the of water in the sample
• 0.75g / 5.25 g 14.3 water

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• Now, lets determine the of water molecules
  trapped in the hydrated salt per formula unit.
• Convert the grams of water to moles.
• 0.75g / 18.01 g/mol 0.0416 moles H2O
• Convert the grams of anhydrous salt to moles.
• 4.50g / 208.24g/mol 0.0216 moles BaCl2
• Divide each by the lesser of the two
• 0.0416/0.0216 1.93 (round to 2)
• 0.0216/0.0216 1
• Write your whole number ratio.
• 1 BaCl2 2 H2O

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• Finally, write your formula for the hydrated salt
• BaCl2 . 2 H2O
• To determine how close you were (or to calculate
  error) calculate your theoretical percent
  water in the above formula as usual
• 2(18.01) / ( 208.24 2(18.01) )
• Mass of water / total mass
• 14.75
• Compare with your original 14.30 experimental
  percent water 14.75 14.30 0.45
• error is this difference divided by the
  theoretical
• 0.45 / 14.75 6.6 error (not too bad for a
  first year chemistry lab student)

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Polyatomic Ions

• As you saw when we discussed and calculated
  oxidation numbers, you sometimes will see a
  cluster of atoms that have a combined overall
  charge
• These are called poly (many) atomic (atoms) ions
  (with a charge)
• Ex. MnO4 -1 the permanganate ion
• I have provided you with a list of polyatomic
  ions that youre responsible for memorizing

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Naming Ionic Compounds

• 1. When naming ionic compounds, begin with the
  cation.
• This can either be a metal from the periodic
  table OR it could be a polyatomic cation
• What are some cations?
• Family I, II, or III, transition metals, or
  ammonium

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Ionic Compounds

• 1.When writing the name, start with the cation
• 2.Follow it by the anion (for now well start
  with polyatomic ions that have a known charge)
• Choose a polyatomic anion that would balance with
  Cu2
• Lets use sulfate for this example
• Write the formula for this compound
• Now write the name of it
• CuSO4

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Using Roman Numerals

• When the cation can vary in its oxidation number/
  charge, we must use a Roman numeral to indicate
  its charge
• Ex. Iron can vary in its ox
• Write the formula for iron bonded to sulfate IF
  the iron has a 3 charge
• Fe2(SO4)3

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Using Roman Numerals

• Fe2(SO4)3
• When we name this compound, we must use a Roman
  numeral to indicate irons charge so the name of
  this compound is
• Iron (III) sulfate because the iron has a 3
  charge
• Do NOT confuse this with the subscript, which
  indicates only how many atoms of iron we have

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Ionic Compounds

• Do we have to use the Roman numeral for family I
  and II elements?
• No, they have a set charge that is understood
• Practice writing the names of the following
• ZnSO4 CaCO3 Fe2(SO4)3
• Zinc sulfate Calcium Carbonate Iron (III)
  Sulfate

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Other Anions

• Polyatomic ions do not all end in ide
• but other ions do
• When naming, you still place the name of the
  cation first
• Followed by the anion (ending in ide)
• What would you call MgO ?
• Magnesium oxide

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Ionic Compounds (cont.)

• Name each of the following
• NaCl KI FrBr
• Sodium chloride Potassium iodide Francium
  bromide
• What happens when you have multiple anions with a
  cation? For example MgCl2 ?
• This is still called Magnesium chloride because
  in order to form this ionic compound, there HAS
  to be 2 chlorines

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