C4 - Chemical Patterns

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  • Atoms are the building blocks of everything.
  • You can only see them with an incredibly powerful electron microscope.
  • Protons are heavy and positively charged.
  • Neutrons are heavy and neutral.
  • Electrons are tiny and negatively charged
  • Neutral atoms have no charge overall.
  • The charge on the electrons, is the same as the change on the protons.
  • The number of protons always equal the number of electrons in a neutral atom.
  • If some electrons are added or removed, the atom becomes charged and is then an ion.
  • The number of neutrons isn't fixed, but usually about the same as the number of protons.
  • The number of protons in an atom decides the element that it is.
  • Elements all have different properties due to differences in their atomic structure.
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The Nucleus and The Electrons

The Nucleus:

  • Is in the middle of the atom.
  • Contains protons and neutrons.
  • Has a positive charge because of the protons.
  • Almost the whole mass of the atom is concentrated in the nucleus.
  • The nucleus is tiny compared to the rest of the atom.

The Electrons:

  • Electrons move around the nucleus.
  • They're negatively charged.
  • They cover a lot of space even though they are tiny.
  • The volume of their orbits determines how big the atom is.
  • The virtually have no mass.
  • They are arranged in shells around the nucleus.
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Chemical Equations

  • During a chemical reaction, nothing appears or dissapears. You have the same atoms, just arranged in different ways.
  • (s) - solid
  • (l) liquid
  • (g) - gas
  • (aq) - dissolved in water
  • A number in front of a formula applies to the whole formula.
  • The little numbers in the middle or at the end of a formula only apply to the atom or brakets immediately before.
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Line Spectrums

  • Lithium (Li) produces a red flame.
  • Sodium (Na) produces a yellow/orange flame.
  • Potassium (K) produces a lilac flame.
  • When heated, the electrons in an atom are excited and they release energy as light.
  • The wavelengths emitted can be recorder as a line spectrum.
  • Different elements emit different wavelengths of light. This is due to each element having a different electron arrangement.
  • Each element has a different pattern of wavelenths, and a different line spectrum.
  • Line spectrums can be used to identify elements.
  • The technique used to produce line spectrums is called sprectroscopy.
  • Spectroscopy has allowed scientists to discover new elements.
  • Caesium and Rubidium were both discovered by their line spectrum.
  • Helium was discovered in the line sprectrum of the sun.
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  • Dobereiner tried to organise elements into trends.
  • In the 1800's, scientists could only measure relative atomic mass, so the known elements were arranged in order of atomic mass.
  • In 1828, Dobereiner started to put this list of elements into groups based on their chemical properties. 
  • He put the elements into groups of three, which he called triads.
  • An example of a triad - Chlorine, Bromine and Iodine.
  • Also, Litium, Sodium and Potassium.
  • The middle element of each triad had a relative atomic mass that was the average of the other two eg.

Lithium has a relative atomic mass of 7.

Sodium has a relative atomic mass of 23.

Potassium has a relative atomic mass of 39.

(7+39) / 2 = 23

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  • Newlands came up with the law of octaves.
  • Newlands notices that when you arranged the elements in order of relative atomic mass, every eighth element had similar properties, so he listed some of the known elements in rows of seven.
  • These sets of eight were called Newlands' Octaves.
  • The pattern failed on the third row, with the transition metals like Titanium (Ti) and Iron (Fe) ruining it.
  • Newlands presented his ideas to the chemical society in 1865. His work was criticised because:-

- His groups contained elements that didn't have similar properties e.g. carbon & titanium.

- He mixed up metals and non-metals e.g. oxygen and iron.

- He didn't leave any gaps for elements that hadn't been discovered yet.

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Dmitri Mendeleev

  • In 1869 in Russia, Mendeleev arranged the known elements into his table of elements, leaving various gaps.
  • Mendeleev put them in order of atomic mass, as Newlands had done before, but he realised he had to leave gaps in order to keep elements with similar properties in the same verticle groups. 
  • The gaps predicted the properties of undiscovered elements.
  • When they were found and they fitted into the pattern, it helped to confirm Mendeleeve's ideas.
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The Modern Periodic Table

  • The Periodic table puts elements with similar properties together. 
  • The elements are laid out in order of increasing proton number.
  • Elements with similar properties form columns called groups. 
  • The group number tells you how many electrons there are in the outer shell.
  • Rows are called periods.
  • You can calculate the number of neutrons by subtracting the number of protons from the relative atomic mass.

The periodic table shows us:

  • The name and symbol of each element.
  • The proton number of each element (showing you how many proton there are in the nucleus, and also how many electrons there are).
  • The relative atomic mass of each element. (showing you the total number of protons and neutrons there are in the nucleus.)
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Electron Shells

  • Electrons always occupy shells (energy levels)
  • The lowest energy levels are filled first.
  • Only a cerain number of electrons are allowed in each shell.
  • 1st shell - 2
  • 2nd shell - 8
  • 3rd shell - 8
  • Atoms prefer it when they have full electrons shells.
  • In most atoms, the outer shell is not full, making the atom want to react.
  • An elements electron arrangement determines its chemical properties.
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Ionic Bonding

  • When atoms lose or gain electrons, they form charged particles called ions.
  • When atoms lose or gain electrons, all they are trying to do is get a full outer shell.
  • When group 1 elements lose an electron, they form positive ions.
  • When group 7 electrons gain an electron, they form negative ions.
  • Oppositely charged ions are attracted to each other, 
  • Compounds formed between group 1 and group 7 elements are held together by ionic bonds.
  • Solid ionic compounds like sodium chloride are made up of a giant lattice of ions. Each latice forms a single crystal.
  • When ionic compound become molten or are dissolved in water, they can conduct electricity. This is because the ions are able to move.
  • The fact that molten compounds of metals and non-metals can conduct electricity, is used as evidence that they are made up of ions.

Reaction of sodium and chlorine:

  • The sodium atoms gives up its outer electron and becomes an Na+ ion. The Chlrorine atom picks up the spare electron and because a Cl- ion.
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Ions and Formulas

  • The charges in an Ionic Compound Add up to zero.
  • Postive Ions - Sodium (Na+), Potassium (K+), Calcium (Ca2+), Iron II (Fe2+) and Iron III (Fe3+).
  • Negative Ions - Chloride (Cl-), Flouride (F-), Bromide (Br-), Carbonate (CO32-) and Sulfate (SO42-).
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Group 1 - The Alkali Metals

  • Group 1 metals include, Lithium, Sodium, Potassium...
  • They all have one outer shell electron.
  • The alkali metals are shiny when freshly cut, but quickly react with oxygen in moist air and tarnish.
  • When Li, Na and K are put in water, they act vigorously and start fizzing furiously whilst moving around the surface. They produce hydrogen. The reaction makes an alkaline solution. A hydroxied of the metal forms. 
  • Alkali metals react vigorously with chlorine. The reaction produces colourless crystal salts.
  • You have to store alkali metals in oil otherwise they would react with the water vapour in the air.

As you go down Group 1, the alkali metals:

  • Become more reactive because the outer shell is more easily lost, because it's further away from the nucleus.
  • They have a higher density because the atoms have more mass.
  • They have a lower melting and a lower boiling point.
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Group 7 - Halogens

  • The group 7 elements include Chlorine, Bromine and Iodine.
  • They all have seven outer electrons. 
  • The halogens form diatomic molecules which are pairs of atoms.
  • As you go down group 7, the halogens become less reactive. They also have a higher melting and boiling point.
  • The halogens are all non-metals with coloured vapours.

Displacement reactions:

  • A displacement reaction is where a more reactive element displaces a less reactive element from a compound. 
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Lab Safety

  • Oxidising - Provides oxygen which allows other materials to burn more furiously. Example: Liquid Oxygen.
  • Highly Flammable - Catches fire easily. Example: Petrol.
  • Toxic - Can cause death either by swallowing, breathing in, or absorption through the skin. Example: Hydrogen Cyanide.
  • Harmful - Like toxic but not as dangerous. Example: Copper Sulfate.
  • Explosive - Can explode. Example: Some Peroxides.
  • Corrosive - Attacks and destroys living tissue, including skin and eyes. Example: Concentrated sulfuric acid.
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Safety Precautions

Alkali Metals:

  • Group 1 elements are really reactive and combust spontaneously.
  • They should never be touched with bare hands.
  • Everything needs to be kept dry so it doesn't react with water.


  • Are harmful. Chlrone and Iodine are both very toxic.
  • Flourine is too dangerous to use in a lab.
  • Must avoid contact with skin as liquid bromine is corrosive.
  • The must be used inside fume cuoboards so you don't breathe in the poisonous vapours.
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