Atoms, Ions and Compounds — Page-by-Page Notes

Page 1

• Topic: Atoms, Ions and Compounds
• Focus areas include Molecular Compounds and Ionic Compounds.
• Source: August 12 Lecture notes (Page 1 content shows core topics and terminology).

Page 2

• Subatomic particles:
  • Protons carry a charge of $+1$ and electrons carry $-1$; neutrons are neutral.
  • Relative masses: protons and neutrons have a relative mass of $1$; the electron has a relative mass of $0$ (mass is so small that we ignore it in this context).
  • Location: protons and neutrons are in the nucleus; electrons orbit the nucleus.

Page 3

• Symbols of Elements:
  • Elements are represented by one- or two-letter symbols (example: carbon is represented by the symbol $ext{C}$).
  • Atomic number: the number of protons in an atom, denoted by $Z$, written as a subscript BEFORE the symbol: $_{Z}^{ }X$ (e.g., for carbon, $_{6}^{ } ext{C}$).
  • Mass number: the total number of protons and neutrons in the nucleus, denoted by $A$, written as a superscript BEFORE the symbol: $^{A}_{Z}X$.

Page 4

In-class #1

• Calculate protons, neutrons, and electrons for:
  • (a) $^{197}_{79} ext{Au}$
  • Protons: $Z = 79$
  • Neutrons: $A - Z = 197 - 79 = 118$
  • Electrons (neutral atom): $Z = 79$
  • (b) Strontium-90 (Sr-90)
  • Strontium has atomic number $Z = 38$
  • Neutrons: $A - Z = 90 - 38 = 52$
  • Protons: $38$
  • Electrons (neutral atom): $38$

Page 5

In-class #2

• Give the chemical symbol including mass-number superscript for:
  • (a) The ion with 22 protons, 26 neutrons, and 19 electrons
  • Z = 22 (Titanium, Ti); N = 26; A = Z + N = 48
  • Electrons = 19 → charge = Z - e = 22 - 19 = +3 (cation)
  • Ion notation: $^{48}_{22} ext{Ti}^{3+}$
  • (b) The ion of sulfur that has 16 neutrons and 18 electrons
  • Z = 16 (Sulfur, S); N = 16; A = 32
  • Electrons = 18 → charge = Z - e = 16 - 18 = -2 (anion)
  • Ion notation: $^{32}_{16} ext{S}^{2-}$

Page 6

Chemical Formulas

• The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound.
• Molecular compounds are composed of molecules.
• Diatomic Molecules: seven elements naturally exist as diatomic molecules, i.e. consist of two atoms each: $ext{H}<em>2, ext{N}</em>2, ext{O}<em>2, ext{F}</em>2, ext{Cl}<em>2, ext{Br}</em>2, ext{I}_2$.

Page 7

Ions

• Cations are formed when at least one electron is lost.
  • Monatomic cations are formed by metals.
• Anions are formed when at least one electron is gained.
  • Monatomic anions are formed by nonmetals.

Page 8

Writing Formulas

• Because compounds are electrically neutral, you can determine the formula as follows:
  • The charge on the cation becomes the subscript on the anion.
  • The charge on the anion becomes the subscript on the cation.
  • If these subscripts are not in the lowest whole-number ratio, divide them by the greatest common factor (GCF).

Page 9

Molar Mass

• Molar mass is the mass of 1 mole of a substance: $M ext{ (g/mol)}$.
• The molar mass of an element is the atomic weight from the periodic table.
• If the element is diatomic, its molar mass is twice the atomic weight of the element.
• The formula mass (in amu) is numerically equal to the molar mass (in g/mol):$M<em>{ ext{formula}} = M</em>{ ext{molar}} ext{ (in amu vs g/mol, respectively)}.$

Page 10

Avogadro’s Number

• In the lab, dealing with individual molecules is impractical; a quantity of $6.02 imes 10^{23}$ atoms or molecules is used to represent a lab-sized amount (ONE MOLE).
• One mole of carbon-12 has a mass of $12.000 ext{ g}$.
• Avogadro’s number: $N_A = 6.02 imes 10^{23} ext{ mol}^{-1}.$

Page 11

Formula Mass and Molecular Mass

• A molecular mass is the sum of the atomic weights of the atoms in a molecule.
• A formula mass is the sum of the atomic weights for the atoms in a chemical formula.
• Ionic compounds use formula weights (not molecular weights).
• Example: Ethane, $ext{C}<em>2 ext{H}</em>6$
  • Molecular mass:
    $M( ext{C}<em>2 ext{H}</em>6) = 2(12.011 ext{ amu}) + 6(1.00794 ext{ amu}) = 30.070 ext{ amu}.$
  • In g/mol:
    $M( ext{C}<em>2 ext{H}</em>6) = 2(12.011 ext{ g/mol}) + 6(1.00794 ext{ g/mol}) = 30.070 ext{ g/mol}.$
• Note: Ionic compounds use formula weights, not molecular weights.

Page 12

In-class #3

• Calculate the molar mass (g/mol) of:

  • (a) Sucrose, $ext{C}<em>{12} ext{H}</em>{22} ext{O}_{11}$
  • (b) Calcium chloride, $ext{CaCl}_2$

• Computations (using standard atomic weights):

  • Sucrose: egin{aligned}M( ext{C}{12} ext{H}{22} ext{O}_{11}) &= 12(12.011) + 22(1.00794) + 11(15.999) \&= 144.132 + 22.17468 + 175.989 \& ext{≈ } 342.29568 ext{ g/mol} \& ext{(≈ } 342.30 ext{ g/mol)}
    ext{.}

  • Calcium chloride: M( ext{CaCl}_2) = 40.078 + 2(35.453) = 40.078 + 70.906 = 110.984 ext{ g/mol} ext{(≈ } 111.0 ext{ g/mol)}.

Page 13

Powers of Ten

• kilo- (k): $10^3$

  • Example: 1 kilometer = $10^3 ext{ m}$

• base unit (meter): $10^0 = 1$

• centi- (c): $10^{-2}$

• milli- (m): $10^{-3}$

  • Examples: 1 meter; 1 centimeter = $10^{-2} ext{ m} = 0.01 ext{ m}$; 1 millimeter = $10^{-3} ext{ m} = 0.001 ext{ m}$

• Questions:

  • How many centimeters are in a meter? $1 ext{ m} = 100 ext{ cm}$
  • How many millimeters are in a meter? $1 ext{ m} = 1000 ext{ mm}$
  • How many meters are in a kilometer? $1 ext{ km} = 1000 ext{ m}$

Page 14

In-class #4

1) Calculate the number of moles of $ext{FeCl}_3$ in $57 ext{ mg}$ of iron(III) chloride.

• Molar mass of $ext{FeCl}_3$:
  • Fe: $55.845 ext{ g/mol}$
  • Cl: $35.453 ext{ g/mol}$ per Cl; $3 imes 35.453 = 106.359 ext{ g/mol}$
  • Total $M( ext{FeCl}_3) = 55.845 + 106.359 = 162.204 ext{ g/mol}$
• Mass given: $57 ext{ mg} = 0.057 ext{ g}$
• Moles: n = rac{m}{M} = rac{0.057}{162.204} \approx 3.51 imes 10^{-4} ext{ mol}