Chapter 2: Atoms, Ions and Compounds

Atom: the smallest part of an element that can't be broken down into anything simpler

Molecule: two or more atoms covalently bonded

Ion: a charged particle, usually formed when atoms lose or gain electrons

Compound: substance containing two or more elements bonded together

Isotope: atoms of the same element with the same number of protons but a different number of neutrons

Electrons determine chemical properties

Isotopes have the same chemical properties because they have the same electron configuration and so react in the same way

Different elements have different chemical properties because they have different electron configurations and so react in different ways

Neutrons determine physical properties

Isotopes have different physical properties because they have a different number of neutrons 

A higher mass means a higher melting / boiling point and higher density

The number of neutrons has no effect on chemical properties / reactions of the element

Atoms:

• Nearly all of an atom's mass is in the nucleus
• The overall charge on a atom is 0
• The number of protons identifies the element

Ions:

• Cations have lost electrons and have fewer electrons than protons so are positive
• Anions have gained electrons and have more electrons than protons so are negative

Heavy water, D2O, contains the H-2 isotope, called deuterium, D

Super heavy water, T2O, contains the H-3 isotope, called tritium, T

The physical properties differ because isotopes have a different number of neutrons

Relative masses: H2O = 18, D2O = 20, T2O = 22

As the relative mass increases, the boiling point, melting point and density also increases

Carbon-12 is the standard for atomic mass measurements

Atomic mass units, u: mass of carbon-12 isotope is defined as exactly 12u

Standard mass = 1u which is 1/12 of the mass of an atom of carbon-12

Mass defect: the small amount of mass lost to hold the nucleus together (binding energy)

Calculated by finding the difference between the mass of the isotope and the mass number of the element

Relative atomic mass:

weighted mean mass of an atom of the element / 1/12 of the mass of an atom of carbon-12

Relative isotopic mass:

mass of an atom of the isotope / 1/12 of the mass of an atom of carbon-12

Calculations:

Ar = (% abundance x mass number) / 100

If relative abundance is used: (relative abundance x mass number) / total relative abundance

• % abundances of isotopes are found using a mass spectrometer
• Ions are detected as a mass-to-charge ratio, m/z
• m/z = relative mass of ion / relative charge on ion

Mass spectrum: abundance on y-axis, m/z on x-axis

Anions                                                       Cations

• OH-           Hydroxide                              Fe2+       Iron
  
• NO3-         Nitrate                                    Zn2+      Zinc
• NO2-         Nitrite                                     Cu2+     Copper
• SO42-       Sulfate                                    Pb2+     Lead
  
• SO32-       Sulfite                                      Ag+       Silver
• CO32-       Carbonate                              NH4+     Amonnium
• PO43-       Phosphate
• HCO3-      Hydrogen carbonate
  
• CN-          Cyanide
• MnO4-      Manganate
• Cr2O72-    Dichromate

Acids

• HCl           Hydrocholoric acid                  H3COOH       Ethanoic acid
• HNO3       Nitric acid                                 HCOOH         Methanoic acid
• H2SO4     Sulfuric aicd
• H3PO4     Phosphoric acid

Diatomic molecules: elements that exist as small molecules, containing only 2 atoms

E.g. H2, N2, O2, F2, Cl2, Br2, I2

The only other elements that exist as small molecules are phosphorus, P4 and sulfur, S8

*In equations, sulfur is written just as S, not S8

Binary compound: only contains 2 elements, ends in IDE

Polyatomic ion: and ion containing atoms of more than one element

Brackets are used when there is more than one polyatomic ion e.g. Mg(OH)2

Formula unit: the formula of an ionic compound