About Formulas

1
About Formulas
2
Types of Bonding
3
Metals vs. Nonmetals

• Most of the elements in the periodic table are
  metals.
• Except for H, everything to the left of the
  staircase is a metal
• Nonmetals are located to the right of the
  staircase.

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Crystal Lattice vs. Molecule

• Molecule Discrete unit, definite number of
  atoms. All molecules of a given compound are
  identical. Exact formula.
• Crystal Lattice Regular, repetitive, 3-D
  arrangement of atoms, ions, or molecules. No set
  size. Not perfect. No two exactly alike.
• Formula gives smallest whole number ratios.
  (Formula unit). Exact formulas NOT useful.

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2-D repetitive patterns
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NaCl
Crystals 3-D repetitive patterns
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Formulas

• Tell the type number of atoms in a compound.
• Microscopic level imagine 1 atom, molecule, or
  formula unit.
• Formula gives atom ratios.
• Macroscopic level imagine working in the lab
• Formula gives mole ratio

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Formulas

• Ex 2H2O can mean
• 2 molecules of water
• 2 moles of water.
• 2H2O molecules have 4 hydrogen atoms 2 oxygen
  atoms.
• 2 Moles of H2O molecules have 4 moles of
  hydrogen atoms 2 moles of oxygen atoms.

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Ionic vs. Covalent Formulas

• By looking at the type of atoms, we can decide if
  the substance is an ionic compound or a molecular
  (covalent) compound.

Molecular compounds usually have all
nonmetals. Ionic compounds usually have metal
nonmetal.
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Empirical Formula

• Smallest whole number ratio of the elements in
  the compound.
• Ionic compounds have only empirical formulas.
• Covalent compounds have empirical and molecular
  formulas. They can be the same or different.

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Molecular Formulas

• For covalent (molecular) compounds.
• Give the exact composition of the molecule.
• Molecular compounds have both empirical and
  molecular formulas. They can be the same or
  different.

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Say everything you can about

• NaCl C6H6 C6H12O6
• CH4 CF4 BeSO4
• C8H18 KI C2H4Cl2
• CaBr2 CuSO4?5H2O

Covalent, molecular
Ionic, empirical
Covalent, molecular
Covalent, empir., molec?
Covalent, empirical, molecular?
Ionic, empirical
Ionic, empirical
Covalent, molecular
Covalent, molecular
Ionic, empirical
Ionic, empirical
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Say everything you can about

• PH3 C2H4 Al2O3
• SrI2 NF3 H2Se
• CH3OH SiO2 H2O2
• CCl4 XeF4 P4O10

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Empirical Formulas from Composition

• Convert to mass.
• Convert to moles.
• Divide by small.
• Multiply til whole.
• Note 1 last step not always used.
• Note 2 sometimes you start at step 2, if they
  give analysis data in grams.

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Compound 63.52 Fe 36.48 S

• Convert to grams Assume 100 g sample-size to
  make math easy.
• 63.52 g Fe and 36.48 g S
• Convert to moles Divide by atomic masses.
• 63.52 g Fe / 55.8 g Fe/mol Fe 1.138 mol Fe
• 36.48 g S / 32.1 g S/mol S 1.136 mol S

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Compound 63.52 Fe 36.48 S

• Divide by small This is where you take the
  ratio.
• 1.136 and 1.138 are essentially the same. Divide
  both by 1.136 and get 1 for each.
• Empirical formula is FeS.

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26.56 K, 35.41 Cr, remainder O

• Assume 100-g sample size.
• 26.56 g K
• 35.41 g Cr
• 38.03 g O
• Convert to moles.
• 26.56 g K / 39.1 g K/mol K 0.679 mol K
• 35.41 g Cr / 52.0 g Cr/mol Cr 0.681 mol Cr
• 38.03 g O / 16.0 g O/mol O 2.38 mol O

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26.56 K, 35.41 Cr, remainder O

• Divide by small
• 0.679 and 0.681 are essentially the same.
• Much smaller than 2.38.
• Divide each by 0.68
• K 1, Cr 1, O 3.5 KCrO3.5? No decs
• Multiply til whole
• K2Cr2O7. Multiply all three by 2, doesnt change
  relationship.

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Relationship between empirical and molecular
formulas

• The molecular formula is a whole number multiple
  of the empirical formula.
• Molec. Formula n (Empirical Form.)
• n is a small whole number, which multiplies the
  subscripts. Sometimes, n 1.

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Molecular Formula

• If you know the empirical formula and the molar
  mass, you can find the molecular formula.
• Step 1 Find the mass of the empirical formula.
• Step 2 Molar mass ? Empirical mass small
  whole number, n
• Step 3 Multiply the subscripts in the empirical
  formula by n.

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Finding Molecular Formulas

• Find the molecular formula for the substance
  whose empirical formula is CH and whose molar
  mass 78.0 g
• Step 1 Empirical mass 13.0 g
• Step 2 Molar mass n 78.0/13.0
• n 6.
• Step 3 6 X subscripts in CH C6H6.

Empirical mass
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What can you say about CuSO4?5H2O?

• Its a hydrated salt.
• For every mole of CuSO4, there are 5 moles of
  water.
• If its heated, it dries out. The water goes
  into the air and youre left with the anhydrous
  salt, CuSO4.
• CuSO4?5H2O ? CuSO4 5H2O

?
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CuSO4?5H2O ? CuSO4 5H2O
?

• The mole ratio in the formula can be used to
  predict how much water would be given off by any
  size sample.
• If you had 2 moles of CuSO4?5H2O, how much water
  would you lose on heating?
• If you had 5 moles of CuSO4?5H2O, how much water
  would you lose?

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Percent Water in CuSO45H2O

• Step 1 Calculate the formula mass.
• Mass of salt Mass of water
• Percent (Part/Whole) X 100
• H2O Mass H2O/Formula Mass X 100

composition from the formula
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Percent Water in CuSO45H2O

• 12.00 grams of CuSO45H2O are heated gently.
• The final mass after repeated heatings is 7.68
  grams. What is the of water?
• Mass of H2O (12.00 7.68) 4.32 g
• Percent H2O 4.32 g X 100 36

12.00 g
Percent composition from experimental data
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CuSO4?xH2O ? CuSO4 xH2O
?

• Use experimental data to find x
• 12.00 grams of CuSO45H2O are heated gently.
  (Hydrated salt.)
• 7.68 g mass anhydrous salt remaining mass
  CuSO4.
• Mass of H2O (12.00 7.68) 4.32 g

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CuSO4?xH2O ? CuSO4 xH2O
?

• Note moles of CuSO4 moles of H2O can be
  determined.
• Have grams already.
• Convert to moles
• Divide by small
• Multiply til whole.

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CuSO4?xH2O ? CuSO4 xH2O
?

• 7.68 g CuSO4 / 159.6 g/mol
• 0.04812 mol CuSO4
• 4.32 g H2O / 18.0 g/mol
• 0.240 mol H2O
• Divide by small, which is 0.04812
• CuSO4?5H2O