Chemistry

• Atomic Structure


  • Atoms are the components that make up all elements
  • Atoms are made up of three types of sub-atomic particles – protons, neutrons, and electrons
  • Protons and neutrons make up the nucleus, where most of the mass is concentrated. Electrons orbit the nucleus in shells
• Subatomic particles


  • Proton
  • Neutron
  • Electron
• Proton


  • Relative Mass: 1
  • Relative Charge: +1
• Neutron


  • Relative Mass: 1
  • Relative Charge: 0
• Electron


  • Relative Mass: 1/1840
  • Relative Charge: -1
• The Evolution of Atomic Structure Over Time

  1. Dalton proposed that all atoms of one element are the same and are different from the atoms of another element. Atoms in his model were tiny and indivisible.

  2. Thomson discovered the electron. He proposed the plum pudding model where negatively charged electrons move in a 'sea' of charge in a positively charged atom.

  3. Rutherford found that most of the mass is concentrated in the positive nucleus, with negative electrons orbiting it. The positive and negative charges balance to make the atom neutral.

  4. Bohr suggested that electrons orbit the nucleus on paths. Bohr's planetary model provided an explanation for the difference in energy of electrons at different distances from the nucleus.

  5. The current model is composed of protons, neutrons, and electrons. Protons and neutrons are found in the nucleus and are made up of smaller quarks, whereas electrons surround the central nucleus.
• Relative Isotopic Mass


  The mass of an atom of an isotope compared with 1/12th of the mass of an atom of carbon-12
• Relative Atomic Mass


  The ratio of the average mass of an atom of an element to 1/12th of the mass of an atom of carbon-12
• Relative Molecular Mass


  The ratio of the average mass of a molecule of an element or compound to 1/12th of the mass of an atom of carbon-12
• Relative Formula Mass


  Similar to relative molecular mass but applies to ionic compounds
• Mass Number (A)


  The total number of protons and neutrons in the nucleus
• Atomic Number (Z)


  The number of protons. The number of positively charged protons is equal to the number of negatively charged electrons in an atom, making the atom neutrally charged
• Mass number = number of protons + number of neutrons
• Atomic number = number of protons = number of electrons
• Ions


  • Formed by atoms losing or gaining electrons
  • A charge of x- means that the number of electrons in the ion is the atomic number + x
  • A charge of x+ means that the number of electrons in the ion is the atomic number - x
• Isotopes


  • Atoms with the same number of protons and different numbers of neutrons. Therefore, they have different mass numbers but the same atomic number.
  • Isotopes of the same element have the same electronic configuration so react in the same way in chemical reactions but have slightly different physical properties.
• Mass Spectrometry


  1. Ionisation - the sample is dissolved in a volatile solvent and ejected through a hollow needle. The needle is connected to a positive terminal of a high voltage supply. This produces tiny positively charged droplets.

  2. Acceleration - ions are accelerated towards a negatively charged plate to give all ions constant kinetic energy. So, the velocity of each ion will depend on its mass.

  3. Ion drift - ions pass through a hole in the negative plate, forming a beam.

  4. Detection - the positive ion picks up an electron which causes a current to flow. Flight times are recorded.

  5. Data analysis - the signal from the detector passes to a computer which generates a mass spectrum.
• Mass Spectrum


  • Gives information about the relative abundance of isotopes on the y axis and about the relative isotopic mass on the x axis.
  • Can be used to determine the relative atomic mass (Ar).
  • For a molecular sample, shows the relative molecular mass on the x axis.
• Electronic Configuration


  • Electrons orbit the central nucleus in shells. Each shell can hold 2n2 electrons, where n is the principal quantum number.
  • Electron shells are made up of atomic orbitals, which are regions in space where electrons may be found.
  • Each shell is composed of one or more orbitals and each orbital can hold one pair of electrons.
  • There are four main types of orbitals: s-, p-, d-, and f-.
  • Electrons have an intrinsic property (spin). For two electrons in the same orbital, the spin must be opposite to minimise the repulsion.
  • Within each shell, orbitals that are of the same energy level are grouped together in sub-shells.
  • There are 1 s-orbital, 3 p-orbitals, 5 d-orbitals and 7-p orbitals possible in each subshell.
  • Sub-shells have different energy levels. Note that 4s is lower in energy than 3d, so 4s will fill first.
  • Shells and sub-shells are filled with electrons according to a set of rules: Atomic orbitals with the same energy fill individually first before pairing, Aufbau principle – the lowest available energy level is filled first, No more than two electrons can fill an atomic orbital.
• Electron Configuration


  Written with n representing principal quantum number. X is the type of orbital and y is the number of electrons in the orbitals of the subshell e.g. potassium has 19 electrons and its electron configuration is written as 1s22s22p63s23p64s1
• Ionisation Energy


  • A measure of the energy required to completely remove an electron from an atom of an element to form an ion.
  • First ionisation energy is the energy required to remove one electron from each atom in one mole of the gaseous element to form one mole of gaseous 1+ ions.
  • Successive ionisation energies apply to the removal of electrons after the first ionisation energy. The nth ionisation energy is: X(n-1)+ (g) → Xn+(g) + e-
• Factors affecting ionisation energies

  • Atomic radii - The larger the atomic radius, the further away the outer electrons are held from the nucleus, and the smaller the nuclear attraction.
  • Nuclear charge - The greater the nuclear charge, the greater the attractive force on the outer electrons.
  • Shielding - electrons repel each other due to their negative charge. The greater the number of inner shells of electrons, the greater the repulsion of the outer shell of electrons.
• Atomic radii and ionisation energy

  The greater the attraction, the harder it is to remove an electron. Therefore, the ionisation energy will be larger.
• Atomic radii show periodicity. Across a period, the radius decreases while down a group, the radius increases.
• Ionisation energy and atomic position

  • Ionisation energy increases across a period, electrons are all added to the same shell resulting in greater attraction.
  • Ionisation energy decreases across a group, as the number of shells increases, so does the atomic radius and shielding, reducing attraction.
• Mole


  • The unit used to quantify the amount of a substance. It can be applied to any amount of chemical species, including atoms, electrons, molecules and ions.
  • The amount of substance that contains the same number of atoms or particles as 12 g of carbon-12.
  • The number of particles in 12g of 12C is the Avogadro constant of 6.022 x 1023 mol-1.
• Concentration


  • The amount of solute present in a known volume of solution.
  • c is the concentration (mol dm-3)
  • n is the number of moles in solution (mol)
  • V is the volume (dm3)
• 1 dm3 = 1000 cm3
• 1 m3 = 1000 dm3
• Molar Gas Volume


  One mole of any gas under standard conditions will occupy the same volume of 24 dm-3 mol-1
• Ideal Gas Equation


  • p is pressure (Pa)
  • V is volume (m3)
  • n is the number of moles (mol)
  • R is the gas constant (8.314 JK-1)
  • T is temperature (K)
• Empirical Formula


  The simplest whole-number ratio of atoms of each element present in a compound.
• Calculating Empirical Formula


  1. Divide the mass of each element by its atomic mass to get the number of moles

  2. Divide each mole value by the smallest to get the whole number ratio
• Molecular Formula


  Gives the number and type of atoms of each element in a molecule. It is made up of a whole number of empirical units.
• Determining Molecular Formula


  Use the empirical formula and relative molecular mass of the molecule
• Balanced Equation


  • No atoms are created or destroyed. The atoms in the reactants rearrange to form the products.
  • There is the same number of atoms of each element in both the reactants and products.
  • State symbols are written after every species to indicate the physical state: Solid (s), Liquid (l), Gaseous (g), Aqueous (aq)
• Ionic Equation


  • Written for any reaction involving ions in solution, where only the reacting ions and the products they form are included.
  • Spectator ions are ions that do not participate in the reaction.
• Empirical formula


  Simplest whole number ratio of atoms of each element in a compound
• Relative molecular mass


  Mass of one molecule of a substance relative to the mass of one atom of carbon-12
• Determining molecular formula from empirical formula and relative molecular mass

  1. Calculate relative molecular mass of empirical formula

  2. Divide relative molecular mass by relative molecular mass of empirical formula

  3. Multiply empirical formula by result to get molecular formula