Chemical Equations and Separation Techniques

• Atoms are the smallest part of an element that can exist.
• Atoms have a radius of around 0.1 nanometres and have no overall charge.
• An element contains only one type of atom.
• Each element is represented by a symbol, e.g., O, Na, or Br.
• Compounds are substances made of two or more elements chemically combined.
• Compounds can only be separated into elements through chemical reactions.
• The central nucleus of an atom contains protons and neutrons.
• Electron shells contain electrons that orbit the nucleus.

Electronic Structures

• The maximum number of electrons that can occupy the first shell is 2.
• The maximum number of electrons that can occupy the second and third shells is 8.
• The maximum number of electrons that can occupy the fourth shells is 2.
• Protons have a relative charge of +1 and a relative mass of 1.
• Neutrons have a relative charge of 0 and a relative mass of 1.
• Electrons have a relative charge of -1 and a very small relative mass.
• The mass number is the sum of protons and neutrons in the nucleus.
• The atomic number represents the number of protons in the atom.
• Number of electrons is equal to the number of protons in an atom.
• Mixtures consist of two or more elements or compounds not chemically combined.
• Mixtures can be separated by physical processes.

Separation Techniques:

• Filtration separates an insoluble solid from a liquid, such as sand from salt water.
• Crystallization separates a solid from a solution, like obtaining pure sodium chloride crystals from salt water.
• Simple distillation separates a solvent from a solution, as in purifying water from salt water.
• Fractional distillation separates a mixture of liquids with different boiling points, such as separating compounds in crude oil.
• Chromatography separates substances that move at different rates through a medium, like separating dyes in food coloring.

Historical Models of the Atom:

• Pre-1900: Atoms were viewed as tiny solid spheres that could not be divided.
• 1897 (Plum pudding): A ball of positive charge with negative electrons embedded in it.
• 1909 (Nuclear model): Features a positively charged nucleus at the center surrounded by negative electrons.
• 1913 (Bohr model): Electrons orbit the nucleus at specific distances.
• James Chadwick provided evidence for the existence of neutrons within the nucleus.

Rutherford’s Scattering Experiment:

• A beam of alpha particles directed at a thin gold foil.
• Most alpha particles passed straight through.
• Some alpha particles were deflected by the positive nucleus.
• A tiny number of particles reflected back from the nucleus.
• Before the discovery of the electron, John Dalton stated that the solid sphere made up the different elements.
• JJ Thomson's experiments showed that an atom must contain small negative charges (discovery of electrons).
• Ernest Rutherford's alpha particle scattering experiment showed that the mass was concentrated at the center of the atom.
• Niels Bohr proposed that electrons orbited in fixed shells; this was supported by experimental observations.
• Law of conservation of mass states total mass of products = the total mass of reactants

Chemical Equations:

• Show chemical reactions.
• Represents reactants and products, and energy changes are often involved.
• Uses words to show reactions.
• Reactants convert into products, e.g., magnesium + oxygen → magnesium oxide.
• Uses symbols to show reactions.
• Reactants convert into products, which need to be balanced.
• E.g., 2Mg + O2 → 2MgO.

Isotopes:

• Atoms of the same element.
• Have the same number of protons but different numbers of neutrons.

Relative Atomic Mass:

• Calculated using the formula: ( (% isotope 1 x mass isotope 1) + (% isotope 2 x mass isotope 2) ) / 100
• Example: for chlorine with isotopes 35Cl (75%) and 37Cl (25%), the relative atomic mass is ((75 x 35) + (25 x 37)) / 100 = 35.5.
• Isotopes that have the same element with the same number of protons and different number of neutrons.

Elements in the Periodic Table:

• Elements are arranged in order of atomic number.
• Elements with similar properties are in columns called groups.
• Elements in the same group have the same number of outer shell electrons, and elements in the same period (row) have the same number of electron shells.
• Early periodic tables were incomplete, and some elements were placed in inappropriate groups.
• This was because the strict order of atomic weights was followed.
• Elements with properties predicted by Mendeleev were discovered and filled missing gaps.
• Knowledge of isotopes explained why order based on atomic weights was not always correct.

Metals and Non-Metals:

• Metals are located to the left of the dividing line in the periodic table.
• Non-metals are located to the right.
• Metals form positive ions.
• They are conductors, have high melting and boiling points, and are ductile and malleable.
• Non-metals form negative ions.
• They are insulators, and have low melting and boiling points.

Group 1 – Alkali Metals:

• Only have one electron in their outer shell and form +1 ions.
• The negative outer electron is further away from the positive nucleus and is easier to lose.
• Very reactive with oxygen, water, and chlorine.
• Reactivity increases down the group.
• With oxygen, they form a metal oxide.
• With water, they form a metal hydroxide and hydrogen.
• With chlorine, they form a metal chloride.

Group 7 – Halogens:

• Have seven electrons in their outer shell and form -1 ions.
• Increasing atomic mass number.
• Increasing proton number means an electron is more easily gained.
• Are diatomic
• Melting and boiling points increase down the group (gas → liquid → solid).
• Reactivity decreases down the group.
  • Metal + halogen → metal halide (e.g., Sodium + chlorine → sodium chloride)
  • Hydrogen + halogen → hydrogen halide (e.g., Hydrogen + bromine → hydrogen bromide)
  • Chlorine + potassium bromide → potassium chloride + bromine (e.g., Cl₂ +2KBr →2KCl + Br2)
• Metal atom loses outer shell electrons and halogen gains an outer shell electron (e.g. NaCl)
• E.g. Cl2 + H2 → 2HCl

Group 0 – Noble Gases:

• Are unreactive and do not form molecules.
• This is due to having full outer shells of electrons.
• Boiling points increase down the group.
• Metal + oxygen → metal oxide
• Metal + water → metal hydroxide + hydrogen
• Metal + chlorine → metal chloride

Transition Metals (Chemistry Only):

• Compared to Group 1 metals, they are less reactive, harder, and denser.
• Have higher melting points.
• Many have different ion possibilities with different charges.
• Used as catalysts.
• Form colored compounds.

Test your knowledge of chemical equations, including word equations and symbolic representations. Explore separation techniques to obtain pure crystals from solutions. Match separation techniques with their applications.