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- Created by: Firdaws123
- Created on: 12-11-17 16:51
Atomic structure
24– Mass number (Protons + Neutrons)
Mg
12– Atomic number (Protons + Electrons)
Relative mass Charge
Proton 1 +1
Neutron 1 0
Electron Very small -1
Atoms are neutral because they have equal number of Protons and electrons so cancel each ther out.
Isotopes
It is an copy of an atom,same number of Protons/electrons but not neutrons.
12 13
C C Reaction same because same number of electrons. Carbon 13
6 6 Will be heavier because it has more neutrons which control mass.
P=6 P=6
E=6 E=6
N=6 N=7
Group 1- Alkali metals
1 electron in outer shell
Called alkali metals because very reactive
React rapidly with oxygen, chlorine and water
As you go down it gets more reactive because......
Outer electron is further from nucleus so less attraction and easier to lose.
Group 0- Noble gases
Full electrons in outer shell
Stable because electrons don’t need to be gained or lost
Very unreactive as full electrons in outer shell
As you go down..
Boiling point increases as relative atomic masses increase
Intermolecular force increases because atoms get bigger which is why force hard to separate
Very unreactive because no electrons moving as full outer shell so no reactions take place
Group 7- Halogens
7 electrons in outer shell
From molecules when when two atoms join by covelent bonds
From covenant compounds when they react with other non-metal atoms
As you go down it gets less reactive because..
Outer shell further from nucleus so less attraction and electron harder to gain
Melting and boiling points increase
Molecules get bigger
Ionic bonding
- When metal and non-metal transfer electrons to become ions
Properties of ionic bonding:
- Giant ionic latitice
- High melting and boiling points so ionic bonds are very strong and lots of energy needed to break them
- Solid ionic compounds don't conduct electricity because no free ions
- Dissolved or melted ionic compounds conduct electricity because ions free to move
-
Metallic bonding
- Positive ions, surrounded by delocalised electrons
- Electrostatic attrcation is attraction between positive object and negative object
- Strong electrostatic attraction between sea of delocalised electrons and positive metal ions
Properties of metallic bonding:
- Giant metallic lattice
- High melting and boiling points so metallic bonds are strong and lots of energy needed to break bonds
- Conducts electricity and heat beacsue free ions. Also electrons carry thermal energy to let metals conduct heat
- Mallebale is shaping ions. When force is applied ions move - Ductile is drawn into wires. When you stretch it, it forms a wire. - In metals layers of atoms are able to slide over each other
Covalent bonding
- Takes place between non-metal elements
- Shaired pair of electrons
- Electroststic attrcation makes covelent bonds very strong
Substances containg covelant bonds have simple molecualr structures:
-The atoms in molecules are held by covelent bonds
-Low melting and boiling points, so gases and liquids
- Don't conduct electricity as no free ions
- Weak intermolecular force between molecules so don't need alot of energy to break
Giant covalent structures
- Always solid at room temperature because they have millions of covelent bonds so high melting and boiling point
Diamond
- Formed from carbon, makes four covelent bonds
- Has loads of covelent bonds which have to be broken when diamond is melted which requires loads of energy. Cannot conduct electricty as no free electrons
Silicon dioxide
-High metling and boiling point because loads of covelent bonds that requires lots of energy to be broken
Graphite Makes 3 covelent bonds. HIgh melting and boiling point as many covelent bonds so lots of energy needed to break bonds. Soft and slippery as no covelent bonds between layers so can slide. Conducts heat and electricity because delocalised electrons that can move.
Graphene and fullerenes
Graphene
-Single layer of graphite,one atom thick. Very strong. Conductor of electricity because delocalised electrons can move. Used in electronics
Fullerenes
-Molecules of carbon atoms with hollow shapes. Hexagonal rings,five or six carbon atoms. Used to deliver drugs into body, for lubricants like machines stopping parts grinding together and catalysts which is to speed up reactiona
Carbon nanotubes
Shaped into long cylinders, rings form 6 carbon atoms, high tesile strength (stretched without breaking) Conductors of heat and electricity
Polymers
- They are plastic
- Solid at room temperature because intermoleculer forces of attrcation are strong so lots of energy needed to break these forces so high melting point.
Conservation of mass
- No atoms are lost or made during chemical reactions so mass of products = mass of the reactants
24g of magnesium racted with 71g of chloride.Calculate mass of magnesium choloride.
Magnesium + chloride ----------> Magnesium cholride
24g + 71g ------------> 95g
The mole
602 000000000000000000000 = 6.02 times 10 to power of 23
Number of moles = Mass in g ------------------------- Mr (Relative formula mass)
You are given a smaple of calcuim carbonate (CaCO3) with a mass of 300g. Calculate number of moles of calcium carbonate in sugar.
CaCo3 40+12+48=100 300 --------- = 3 moles 100
Concentration
Concentration (mol/dm3)= Number of moles ----------------------- Volume (dm3)
If volume in cm-3 divide by 1000.
A solution has a concentration of 0.5 mol/dm3. Calculate the number of moles in 0.2dm3.
0.5 times 0.2 = 0.1 moles
Limiting reagents and excess
Home many moles of zinc iodine would be produced if we used 0.5 moles of zinc and 1 mole of iodine?
Zinc + Iodine-----------> Zinz iodine
1 mole + 1 mole-----------> 1 mole
0.5 + 1 mole------------> 0.5 moles
Limiting + Excess
Redox
Oxidation and reduction
Oxidation
-Gain of oxygen
-Loss of electrons
Reduction
-Loss of oxygen
-Gain of electrons
Reactions of metals
Metals and Oxygen
When a metal recats with oxygen, it forms metal oxide
Magnesium + oxygen ------------> Magnesium oxide
Metals ands acids
All acids contain hydrogen. Hydrochloric ----> HCl Sulfuric ---->H SO Nitric---->HNO When acids react with metals they form salt and hydrogen gas.
Acid + metal --------> Salt + Hydrogen Hydrochloric acid + magnesium -----> magnesium chloride + Hydrogen
Magnesium is more reactive than hydrogen so magnesium can easily displace hydrogen from acids.
Metals and water
When metals react with water a metal hydroxide and hydrogen has formed.
Metals
Unreactive metals
Very unreactive metals are found in ground E.g gold,silver,platinum
Ores
Metals found in ores have reacted with other elements. Oxygen and other impurities(sand). Ores are rocks containing enough metal to make it cost effective to extract.
Smelting iron
1) Iron is placed in a blast furnace. Limestone and carbon is added.
2) Carbon displaces iron to form iron and carbon dioxide. Carbon takes care of oxygen +limestone takes care of impurities.
You can use carbon to extract any metal less reactive than carbon. Iron oxide is heated with carbon. Carbon is more reactive than iron so displaces the metal.
Smelting copper, same extraction as iron
Phytomining + bio-leaching
Phytomining
1) Plants absorb copper, making copper
2) Burn plants and get copper from ash
3)Use electricity (electrolysis) to get metal
Bioleaching
1) Microorganisms absorb copper
2) Use electricity (electrolysis) to get metal
Iron vs Steel
-Both made from iron
Iron
When iron is extracted it is almost pure. Pure iron is soft because atoms are in layers and can slide over each other. Pure iron is brittle (breaks easily) because impurities such as carbon.
Steel
Alloy of iron combined with carbon. Low carbon steel has small amount of carbon and can be shaped easily. High carbon steel harder than low carbon steel. Lots of carbon atoms distort layers and prevent them from sliding.
More reactive metals
Electrolysis is used for more reactive metals. Using electricity is very expensive. Aluminium is extracted with electrolysis because carbon can't displace it as aluminum is more reactive then carbon.
Acids,Alkalis and bases
Acids produce hydrogen ions ( H ) in aqeus solutions. ( Dissolve in water).
HCl (aq) ------>H (aq) + Cl (aq)
Bases are chemicals which can neutralise acids and produce salt and water. Some bases are solid, if solid dissolves in water it is an alkali.
acid + base ----------> salt +water H (aq) + OH (aq) -------> H O
Aqueus solutions of alkalis contain hydroxide ions OH
NaOH (aq) ---------> Na(aq) + OH (aq)
How to know if something is acid or alkali?
-PH probe, determines PH electronically. Tells you number of PH.
-Universal indicator, changes colour whether it is acid or alkali. Red shows acid. Blue shows Alkali.Green shows neutral.
-Litmus paper, red litmus paper only works with alkali and blue with acid.
Strong and weak acids
Strong acids fullt ionise in aqueous solutions Positive and negative part seperates completely.
HCl ------>H + Cl -Hydrochloric acid(stomach acid) -Sulferic acid -Nitric acid
Weak acids partially ionise in aqeuous solutions. H CO (aq) Reverse H (aq) + HCO -Citric acid(lemons) -Ethanic acid (Vinegar) -Phosperic acid (fizzy drinks) -Carbonic acid
The PH scale
A high concentration of H ions (acids) will have low PH.
A low concentration of H ions (basic) will have high PH.
As the PH scale decreases by 1 unit, concentration of hydrogen ions increases by 10 times.
Making salts
Salts produced are sometimes soluble and insoluble.
If you have 2 types of alkali liquids and don't know if it is a metal or metal carbonate. Put acid in and if it bubbles you know gas is produced which is what metal carbonate produces.
5 ways to make salts
1. Neutralisation: Acid + metal hydroxide Alkali are soluble bases. Acid + Alkali ------------> Salt + water
2.Neutralisation: Acid + Metal oxide Bases are insoluble Acid + base --------------> Salt + water
3.Neutralistaion: Acid + Metal carbonate Metal carbonates are insoluble. Acid + base -------------> Salt + water + carbon dioxide
4.Acid + Metal ----------> Salt + hydrogen ( Dangerous to use this)
5. Ammonia NH .Acts as a base Can react with acids to make ammonium,salts which are used as fertiliser.
Rection
1) Reaction of acids with metals
Magnesium + hydrochloric acid -------> Magnesium chloride + hydrogen Mg + 2HCl ---------> MgCl + H
2) Reaction of acids with metal hydroxides
Sodium hydroxide + Nitric acid ----------->Sodium nitrate + Water NaOH + HNO -----------> NaNO + H O
3)Reaction of acids with metal oxides
Calcium oxide + Nitric acid ---------------> Calcium nitrate + Water CaO + 2HNO ---------------> Ca (NO ) + H O
4) Reaction of acids with metal carbonates
Sodium carbonate+ Nitric acid---------------> Sodium nitrate+water + carbon dioxide Na CO + 2HNO --------------> 2NaNO + H O + CO
Electrolysis
Splitting of ionic compounds using electricity. ionic compounds contain metals combined with non-metals.E.g Sodium (metal) Chloride(non-metal) (NaCl)
Ionic compounds are made up of positive + negative ions. As solid ionic compounds cannot conduct electricity because ions cannot move around. But when ionic compounds are melted, the ions are free to move and conduct electricity. These liquids are called electrolysis.
Negative ions are attracted to positive elctrode ( Anode).
Positive ionss are attracted to negative electrode (Cathode)
Electrolysis of copper chloride
Chloride is attracted to the positive electrode because chloride is negative. At the positive electrode chloride loses electrons and turn into chlorine gas. Sodium does not form solids because it's more reactive than hydrogen, instead hydrogen gas is made.
Exothermic and endothermic
Exothermic reactions transfer energy to the surroundings. Often a temperature change. E.g When bonfire burns, it transfers heat energy to the surroundings.Objects near bonfire become warmer.
Examples:
-Burning fuels
-Neutralistaion- If beaker is warm
- Resperation- Cells (Releases heat due to our body temp)
Endothermic take in energy from the surroundings. Energy usually transferred as heat energy, causing the reaction mixture and it's suroundings to get colder.
Reversible reactions
Normal reaction
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