C4: atoms elements and the periodic table

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C4: atoms elements and the periodic table
1 HISTORY OF THE PERIODIC TABLE-
1.1 Modern Periodic Table: based on DMITRI MENDELEEV a Russian scientist ideas. He arranged the element into groups (vertical columns) and periods (horizontal rows) based on their relative atomic masses and patterns in properties
1.1.1 Mendeleev left gaps for undiscovered elements and predicted properties of missing elements
1.1.1.1 when the elements were discovered they fit his gaps perfectly
1.1.1.1.1 data about properties of elements in the Periodic Table can be used to work out trends and to make predictions.
1.2 JOHANN DOBEREINER- came up with the TRIAD arrangement that linked patterns of the relative atomic masses for three elements
1.2.1 scientists rejected the traids theory as most elements did not fit their patterns
1.3 JOHN NEWLANDS- came up with the octaves pattern where every eighth element had similar properties
1.3.1 scientists rejected the octaves theory because most elements did not fit their patterns
2 INSIDE THE ATOM
2.1 an element contains all the same type of atoms
2.1.1 all the atoms of a particular element have the same number of protons however an elements atoms can have different isotopes- so they have different numbers of neutrons
2.2 atoms have a tiny, central nucleus that contains protons and neutrons
2.3 the electrons travel around the outside of the atom in shells
2.4 Proton- charge of +1 and a mass of 1
2.4.1 Electron- charge of -1 and a negligible mass
2.4.1.1 Neutron- 0 charge but mass of 1
3 LINES OF DISCOVERY
3.1 when elements are heated they emit coloured flames
3.1.1 some elements emit distinctive flame colours
3.1.1.1 e.g. litium salts produce a red flame
3.2 the coloured light can be split into a line spectrum
3.2.1 the line spectrum is unique to each element
3.3 spectroscopy led to the discovery of some elements
3.3.1 e.g. helium was discovered when chemists looked at the line spectrum from the sun
3.3.2 sodiums line spectrum is two single yellow lines so close together that they look like one
4 PERIODIC TABLE
4.1 is number of protons + number of neutrons= relative atomic mass
4.2 elements are arranged in order of proton number
4.3 Left-->Right along a period = metals to non metals
4.3.1 Left-->Right along a period= each element has one more proton and one more electron
4.3.1.1 number of electrons in outer shell increases across a period
4.3.1.1.1 group number represents number of electrons in outer shell
4.4 elements in Group 0 have a full outer shell so they are INERT- very unreactive
4.5 atoms of elements with up to 3 electrons in their outer shells are metals
4.6 atoms of elements with 5 or more electrons in their outer shells are non metals
4.7 properties e.g. melting points change across a period. These changed are called TRENDS
5 SORTING ELECTRONS
5.1 electrons are arranged in shells around the nucleus
5.1.1 the first shell is closest to the nucleus and can hold up to two electrons
5.1.1.1 the second and third shells can hold 8 electrons
5.1.1.1.1 electron arrangement can be written as a ratio of sorts
5.1.1.1.1.1 . oxygen can be written as 2.6 and chlorine can be written as 2.8.7
5.1.1.1.2 for the first 20 elements the third shell holds 8 electrons
5.1.1.2 each shell holds different energy levels but the closer the electron shell to the nucleus the lower the energy level
6 GROUP 1 THE ALKALI METALS
6.1 all elements in group one have one electron in their outer shell
6.1.1 the physical properties e.g. melting and boiling point decrease down the group. they show a trend down the group and therefore reactivity increases down the group.
6.1.1.1 the outer shell is easiest to lose if they atom is bigger because the electron is further away from the nucleus so reactivity increases down group 1 as the atoms get bigger
6.1.1.1.1 REACTIONS OF G1 WITH CHLORINE
6.1.1.1.1.1 sodium reacts V vigorously to give a yellow flame and white solid sodium chloride
6.1.1.1.1.2 the other G1 react similarly however reaction time gets faster down thee group
6.1.1.1.1.3 metal+chlorine--> metal chloride
6.1.1.1.2 REACTIONS OF G1 WITH WATER
6.1.1.1.2.1 lithium and sodium fizz and move around the surface of the water. The reaction gets more vigorous as you go down the group
6.1.1.1.2.1.1 e.g. potassium explodes and Rubidium reacts even more violently
6.1.1.1.2.2 in the reaction hydrogen gas is formed (squeaky pop) and a metal hydroxide which turns the indicator blue
6.1.1.1.2.2.1 the group 1 metals hydroxides are harmful and corrosive they are very flammable and should be kept away from water and naked flames
6.1.1.1.2.3 METAL+ WATER-->METAL HYDROXIDE + HYDROGEN
6.1.1.1.2.3.1 2NA (s)+ 2H20 (l)-->2NAOH (aq) + H2 (g)
6.2 they are all metals
6.2.1 when they become ions they lose their only electron in their outer shell and become positively charged ions
6.3 they are soft and can be cut with a knife, the freshly cut surface is shiny but it tarnishes quickly in moist air reacting with oxygen
6.4 LI, Na, K, Rb, Cc, Fr
7 GROUP 7 THE HALOGENS
7.1 all have 7 electrons in their outer shell
7.2 F, Cl, Br, I, At
7.3 all diatomic so they have two atoms per molecule
7.4 appearance at room temperature-
7.4.1 Chlorine- pale green gas
7.4.2 Bromine- reddish brown liquid but as a gas it is a reddy brown colour
7.4.3 Iodine- dark grey solid but as a gas its purple
7.5 the physical properties of the alkali metals show a trend down the group. e.g. their melting and boiling points increase
7.5.1 so the reactivity decreases down the group because its even harder to remove electrons/ change them
7.5.1.1 the halogens are less reactive down the group e.g. iron reacts vigorously within contact of fluorine but only reacts with iodine when its heated
7.5.1.1.1 sodium reacts vigorously with chlorine but less violently with iodine
7.6 they are all corrosive and toxic and must be used in a fume cupboard
7.7 REACTIONS WITH METALS
7.7.1 G7 react with G1 and other metals to form metal halides
7.7.1.1 iron+ chlorine--> iron chloride
7.7.1.2 sodium+bromine--> sodium bromide
7.7.1.3 potassium+ iodine--> potassium iodide
7.8 DISPLACEMENT REACTIONS
7.8.1 happens when a more reactive halogen takes the place of a less reactive halogen in a compound
7.8.1.1 chlorine is more reactive than bromine so it will displace bromine in a solution of potassium bromide, but bromine is not as reactive as chlorine so it cannot displace it
7.9 the trend in reactivity is linked to the number of electron shells in the atom. For non metals the smaller the atom (the fewer the electron shells) the more reactive the element
8 IONIC COMPOUNDS
8.1 compounds of group 1 elements and group 7 elements are solids with very high melting points e.g. sodium chloride
8.1.1 these compounds are ionic compounds because they contain charged particles, or ions, that are arranged in a regular pattern called a crystal lattice
8.2 they conduct electricity when molten or in solution of water
8.2.1 this is because the ions are free to move and can move to the electrodes to complete the circuit
8.3 an atom of G1 loses its one outer shell electron to become a positive ion e.g. Na --> Na+
8.3.1 an atom of G7 gains an electron to complete its outer shell and it becomes a negative ion. e.g. Cl--> Cl-
8.3.1.1 the ions then have the same charge as atoms in G0 e.g. Na:: 2.8.1--> Na+ 2.8. which is the same electron arrangement as neon
8.3.1.1.1 Cl: 2.8.7--> 2.8.8 which is the same electron arrangement as argon
8.4 in the formula for an ionic compound the number of positive charges and negative charges balance e.g. Na+ and Cl- become NaCl because the singular positive and negative charges balance each other out. They form the compound NaCl which is sodium chloride
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