chapter 2 atomic structure and isotopes

• protons have virtually the same mass as a neutron
• An electron has a negligible mass of 1/1836th the mass of a proton

• A proton has a charge of +1
• An electron has a charge of -1
• A neutron has no charge

• Nearly all the atoms mass is in the nucleus
• Atoms contain the same number of protons as neutrons
• The overall charge is 0
• Neutrons can be thought of as the glue that holds the nucleus together
• Atoms contain the same number of or slightly more neutrons as protons
• As the nucleus gets larger more and more neutrons are needed

• The number of protons in an atom defines the element
• The atomic number is the number of protons in an element

• Unlike protons, the number of neutrons in the atoms of an element can change
• Isotopes are atoms of the same element with different numbers of neutrons

• Different isotopes of the same element will have the same number of electrons
• Because the number of neutrons has no effect on the reaction, different isotopes will react in the same way
• There may be small changes in physical properties

• Cations (positive ions) have fewer electrons than protons
• Anions (negative ions) have more electrons than protons
• Ions and atoms of the same element will have the same number of protons but a different number of electrons

• Mass defect is the small amount of mass lost due to the strong nuclear force holding together protons and neutrons
• A standard isotope is needed to base all atomic masses
• The mass of a carbon-12 isotope is defined as exactly 12 atomic mass units
• The standard mass for atomic mass is 1u which is equal to 1/12th of an atom of carbon-12
• On this scale, 1u is equal to the mass of a proton or neutron

• Relative isotopic mass is the weighted mean mass of an atom relative to 1/12th of an atom of carbon-12

• Relative atomic mass is the weighted mean mass of an atom relative to 1/12th of an atom of carbon-12
• The weighted mean mass takes account of the percentage abundance and relative isotopic mass of each isotope

• A sample is placed in a mass spectrometer 
• The sample is vaporised and then ionised to form positive ions 
• The ions are accelerated
• Heavier ions move more slowly and are more difficult to deflect than lighter ions
• Ions of each isotope are separated
• The ions are detected on a mass spectrum as a mass-to-charge ratio

mass to charge ratio = relative mass of ion / relative charge on ion

• Metals lose electrons to form positive ions
• Non-metals gain electrons to form negative ions
• Some metals such as some of the transition metals can form several ions with different charges

• A compound that contains two elements only

• An ion that contains atoms of more than one element

• The overall charge must be zero, so the sum of the positive charges must equal the sum of the negative charges

• The number of each ion is shown in subscript after the ion 
• The ionic charge is usually omitted in a full formula
• Brackets are used if there is more than one polyatomic ion

• Elements are shown simply as their symbol unless they exist as a diatomic molecules

• Covalent compounds do not contain ions
• For ionic compounds the formula worked out from the ionic charges is used in the equation- this is called the formula unit

• (g) - gas
• (l) - liquid
• (s) - solid
• (aq) - aqueous

• You must not change any chemical formula
• Balancing numbers go in front of chemical formula
• The equation is balanced when the same number of atoms are on both sides