Module 2: Foundations in Chemistry
Chapter 2: Atoms, ions and compounds
2.1 Atomic structure and isotopes
Properties of protons, neutrons and electrons
Mass
Atoms and their subatomic particles have tiny mass
Chemists use relative masses
o Proton and neutron virtually same mass (neutron slightly greater = 1.001375)
o Electron has around 1/1836th the mass of a proton
Building the atom
Nearly all atom’s mass is in the nucleus
Neutron can be thought as holding the nucleus together, despite the electrostatic
repulsion between its positively charge proton
Most atoms contain the same number/slightly more neutron than protons
As nucleus gets larger, more neutrons are needed
Atomic number
No. of proton in an atoms defines the element
Atomic number = proton number = Z
Mass number = nucleon number = A
Isotopes
Isotopes = atoms of the same element with different number of neutrons and different
masses
Most elements are made up of a mixture of isotopes
Chemical reactions involve the electrons which means that different isotopes of an
element react in the same way
o They have the same no. of electrons, just a different number of no. of neutrons
which has no effect on reactions.
Maybe a small difference in physical properties
o Higher mass isotopes having higher melting point, boiling point and density.
Heavy water
Used to control processes in nuclear reactor
All molecules of H2O contain the deuterium (1 neutron), often referred with its own
symbol D → D2O
Properties of water and heavy water
Physical properties Normal water Heavy water
Melting Point / ⁰C 0.00 3.80
Boiling Point / ⁰C 100.00 101.40
Density / gcm3 1.00 1.11
Atomic structure of ions
Ion is a charged atom, meaning no. of electrons is different from no. of proton
Ions and atoms of an element have the same no.proton but different no.electron
o Positive ions - cations, are atoms with fewer electrons than proton
, o Negative ions - anions, are atoms with more electrons than proton
2.2 Relative mass
Carbon-12
To find the relative mass of an isotope might seem sensible to add together the relative
masses of the 3 subatomic particles. nonono
Strong nuclear force holding together the protons and neutrons comes at the expense of
the loss of a fraction of their mass - mass defect
Carbon-12 - international standard for the measurement of atomic masses
One atom of carbon-12 - 1.992645538 x 10^-26
New unit called the atomic mass unit u is used
o Mass of a carbon-12 isotope is defined as 12u
o Standard mass for atomic mass is 1u, the mass of 1/12th of an atom of carbon-
12
o On this scale, 1u is approximately the mass of a proton or neutron
Relative isotopic mass
Relative isotopic mass - The mass of an isotope compared with 1/12th of the mass of an
atom of carbon-12
No unit since it is a ratio of 2 masses
In most cases, we can assume that the relative isotopic mass is the same as the mass
number A of the isotope (no. protons and no. neutrons)
Relative atomic mass
Relative atomic mass (Ar) - the weighted mean mass of an atom of an element relative
to 1/12th of the mass of an atom of carbon-12
Most elements contain a mixture of isotope, with different relative isotopic mass
Weighted mean mass take account
o Percentage abundance
o Relative isotopic mass
Percentage abundances of the isotopes in a sample of an element are found
experimentally using a mass spectrometer
o Sample is placed in the mass spectrometer
o Sample is vaporise and then ionised to form positive ions
o Ions are accelerated. Heavier ions move more slowly and are more difficult to
deflect than lighter ions → ions are separated
o Ions are detected on a mass spectrum as a mass-to-charge ratio m/z. Each ion
reaching the detector adds to the signal, so the greater the abundance, the larger
the signal.
o Mass-to-charge ratio m/z = relative mass of ion / relative charge on ion
o Relative atomic mass = contribution of isotopes (abundance x mass) / 100
2.3 Formulae and equations
Simple ions from the periodic table
Atoms of metal lose electrons to form cations
Atoms of non-metal gain electrons to form anions
Some transition metals can form several ions with different charges, which is shown with
a roman numeral
o Copper (Ⅰ) - Cu+, copper (Ⅱ) - Cu 2+
o Iron (Ⅱ) - Fe 2+, Iron (Ⅲ) - Fe3+
Chapter 2: Atoms, ions and compounds
2.1 Atomic structure and isotopes
Properties of protons, neutrons and electrons
Mass
Atoms and their subatomic particles have tiny mass
Chemists use relative masses
o Proton and neutron virtually same mass (neutron slightly greater = 1.001375)
o Electron has around 1/1836th the mass of a proton
Building the atom
Nearly all atom’s mass is in the nucleus
Neutron can be thought as holding the nucleus together, despite the electrostatic
repulsion between its positively charge proton
Most atoms contain the same number/slightly more neutron than protons
As nucleus gets larger, more neutrons are needed
Atomic number
No. of proton in an atoms defines the element
Atomic number = proton number = Z
Mass number = nucleon number = A
Isotopes
Isotopes = atoms of the same element with different number of neutrons and different
masses
Most elements are made up of a mixture of isotopes
Chemical reactions involve the electrons which means that different isotopes of an
element react in the same way
o They have the same no. of electrons, just a different number of no. of neutrons
which has no effect on reactions.
Maybe a small difference in physical properties
o Higher mass isotopes having higher melting point, boiling point and density.
Heavy water
Used to control processes in nuclear reactor
All molecules of H2O contain the deuterium (1 neutron), often referred with its own
symbol D → D2O
Properties of water and heavy water
Physical properties Normal water Heavy water
Melting Point / ⁰C 0.00 3.80
Boiling Point / ⁰C 100.00 101.40
Density / gcm3 1.00 1.11
Atomic structure of ions
Ion is a charged atom, meaning no. of electrons is different from no. of proton
Ions and atoms of an element have the same no.proton but different no.electron
o Positive ions - cations, are atoms with fewer electrons than proton
, o Negative ions - anions, are atoms with more electrons than proton
2.2 Relative mass
Carbon-12
To find the relative mass of an isotope might seem sensible to add together the relative
masses of the 3 subatomic particles. nonono
Strong nuclear force holding together the protons and neutrons comes at the expense of
the loss of a fraction of their mass - mass defect
Carbon-12 - international standard for the measurement of atomic masses
One atom of carbon-12 - 1.992645538 x 10^-26
New unit called the atomic mass unit u is used
o Mass of a carbon-12 isotope is defined as 12u
o Standard mass for atomic mass is 1u, the mass of 1/12th of an atom of carbon-
12
o On this scale, 1u is approximately the mass of a proton or neutron
Relative isotopic mass
Relative isotopic mass - The mass of an isotope compared with 1/12th of the mass of an
atom of carbon-12
No unit since it is a ratio of 2 masses
In most cases, we can assume that the relative isotopic mass is the same as the mass
number A of the isotope (no. protons and no. neutrons)
Relative atomic mass
Relative atomic mass (Ar) - the weighted mean mass of an atom of an element relative
to 1/12th of the mass of an atom of carbon-12
Most elements contain a mixture of isotope, with different relative isotopic mass
Weighted mean mass take account
o Percentage abundance
o Relative isotopic mass
Percentage abundances of the isotopes in a sample of an element are found
experimentally using a mass spectrometer
o Sample is placed in the mass spectrometer
o Sample is vaporise and then ionised to form positive ions
o Ions are accelerated. Heavier ions move more slowly and are more difficult to
deflect than lighter ions → ions are separated
o Ions are detected on a mass spectrum as a mass-to-charge ratio m/z. Each ion
reaching the detector adds to the signal, so the greater the abundance, the larger
the signal.
o Mass-to-charge ratio m/z = relative mass of ion / relative charge on ion
o Relative atomic mass = contribution of isotopes (abundance x mass) / 100
2.3 Formulae and equations
Simple ions from the periodic table
Atoms of metal lose electrons to form cations
Atoms of non-metal gain electrons to form anions
Some transition metals can form several ions with different charges, which is shown with
a roman numeral
o Copper (Ⅰ) - Cu+, copper (Ⅱ) - Cu 2+
o Iron (Ⅱ) - Fe 2+, Iron (Ⅲ) - Fe3+
