1. Test Tube 2. Boiling Tube 3. Beaker
4. Conical Flask 5. Test Tube Holder/ Tongs /Clamp Stand
6. Measuring/Graduating 7. Funnel 8. Pipette Cylinder
9. Spatula 10. Tripod 11. Bunsen Burner
12. Glass Bulb 13. Burette 16. Condenser
Changes of State
Filtering Solutions 03-Nov
The Atomic Structure
Atoms contain electrons which travel along electron shells, surrounding the nucleus that contains nucleons such as protons and neutrons. Electrons, neutrons and protons all make up an atom, they are called sub-atomic particles.
Protons – has a relative charge of 1+ and a relative atomic mass of 1
Neutrons – are neutral, therefore the charge is 0. Also has a R.A.M of 1
Electrons – has a negative charge (1- ), and has a R.A.M of 1/2000
Listed in the periodic table, are elements. All elements have a neutral charge, or the charge always equal 0. This is why the amount of protons and electrons are the same in every element. Also, the relative atomic mass does not include the electrons, because their mass is so small, it barely makes up the atom.
Atomic Number & R.A.M To find the numbers of protons, neutrons and electrons of an element, we use the atomic number and the relative atomic mass (R.A.M.) to work it out.
Therefore: Protons = 3 Electrons = 3 Neutrons = (7-3) 4
Electron Structures 04-Nov
Number of electrons allowed of main electron shells:
1st Shell – 2 electrons
2nd Shell – 8 electrons
3rd Shell – 8 electrons
4th Shell – 2 electrons
When illustrating electron structures, electrons should always be filled from the innermost electron shell first (closest to the nucleus), moving outwards.
Electrons are arranged in shells around the nucleus of an atom, the innermost electron shell, having the lowest energy level, is filled first with a maximum of 2 atoms. Also, electron structures do not require you to draw the protons and neutrons of the nucleus; instead the element symbol is placed.
The 1st main shell can be filled up to 2 electrons, the 2nd with 8 electrons, the 3rd with 8 electrons, and the 4th with 2 electrons. When all shells are filled with 2, 8, 8, 2 all electron shells are complete.
Calculating the relative atomic mass
Magnesium + Oxygen ( Magnesium Oxide
Mg + O ( MgO
2Mg (s) + O (g) ( 2MgO (s)
The Periodic table is in order of the proton number or atomic number, a pattern which was designed and discovered by a Russian Scientist, Dmitri Mendeleev. The table helps make sense of the different properties of elements and their ions and help predict how they will behave in different situations. The atomic number shows the number of protons in the atom of and element, and the relative atomic mass tells us the comparative measurement of mass of one atom or number of neutrons and protons together of an atom.
Groups & Periods The numbers listed at the top of the table are called Groups; they show how many electrons are present at the outermost electron shell. There are 8 groups in the table. A group is all the vertical columns of elements, and all the elements in each group have similar properties, they are a chemical family. Periods are the rows going horizontally across the Periodic table, the first period has 2 elements, and the second one has 8, and so on.
Metals & Non-metals In the table above, the red squares of elements represent metals and the yellows are non-metals. The green squares of elements which create a border line between metals and non-metals are called semi-metals or metalloids. Metalloids behave as a metal in some ways but like a non -metal in others. For example, Silicon is shiny like a metal, but brittle like a non-metal. Tin-oxide also reacts both like a metal and non-metal oxide. As this behaviour is like an acid and a base, it is called an amphoteric oxide.
Properties of metals: Shiny, magnetic, conductor of heat, conductor of electricity, strong, ductile, malleable, high-melting points, dense, sonorous.
Properties of non-metals: dull, poor conductor of heat, poor conductor of electricity, low-melting points, brittle.
Metals Metals are located to the left of the borderline of the periodic table. They are able to conduct heat and electricity. They are called conductors. They are able to conduct electricity because they have delocalized electrons which can easily enter and exit the atom. The looser the ‘sea’ of electrons is in a metallic structure, the less resistance and the better conductor.
Reactivity and Forces of Attraction To react, an element needs to complete all their main electron shells. In group 1, all metals have one electron on their last shell; therefore it would be easier to complete the shells by taking away that one electron. When this happens, there is an unbalanced charge as the numbers of positive protons are not the same as the negative electrons; thus the atom becomes a positive ion.
Going down group 1 elements of the Periodic table, they increase in reactivity. This is because the atomic number increases going down. This tells us that the number of protons thus electrons is increasing, meaning more electron shells are needed. When there are electron shells in between the nucleus and the outer electrons, the shielding effect commences. This weakens the force of attraction between the positive protons and negative outer electron of atoms. The distance created by the shells between the protons and outer electron also weakens the atom’s force of attraction. Therefore it is easier to loose that one outer electron during a reaction; and the element becomes more reactive than others. However, if the force of attraction is stronger, it becomes more difficult to separate the outer electron from the atom, and the element is unreactive in comparison to the others. Conversely, in group 7, it is easier to add one more electron rather than taking away 7 electrons to complete the shells. Therefore where the force of attraction is stronger (top of the group) the easier it is to attract an extra electron to the atom. Henceforth the element is more reactive. As a result, group 1 elements increase in reactivity going down the Periodic table, but group 7 elements decrease in reactivity going downwards.
Ions and Ionic Compounds Group 1 ions have a charge of 1+. This is because electrons have a negative charge and protons have a positive charge. When the one electron is taken away and the elements are left with one extra positive proton, the charge which was neutral in the beginning (as there were the same number of protons as electrons) becomes unbalanced. In group 7 ions, the charge is 1- because another negative electron is added to the atom. Thus there is an extra negative charge when compared to positive charges.
REDOX is reduction oxidation.
Oxidation means losing electrons and reduction means gaining electrons. An easy way to remember this is through OIL RIG. Oxidation is loss, reduction is gain. As an example, here is a reaction between iron and copper (II) sulfate:
Fe(s) + Cu2+ (aq) [pic] Fe2+ (aq) + Cu(s)
Iron metal (Fe(s)) loses 2 electrons to form iron ions (Fe2+ (aq))
Fe(s) - 2e-(aq) [pic] Fe2+ (aq)
This is called oxidation. Iron is oxidized.
Copper ions (Cu2+ (aq)) gain 2 electrons to form copper metal (Cu(s))
Cu2+ (aq) + 2e-(aq) [pic] Cu(s)
This is called reduction. Copper is said to be reduced. Oxidation and reduction always occur together. They are called redox reactions.
Oxidation may also be defined as gaining oxygen and reduction defined as losing oxygen.
Group 1 – Alkali Metals 15-Nov The elements in group 1 are all alkali metals. When reacting with water the solution becomes alkaline. Group 1 metals also don’t have many uses as elements; this is because they are too reactive. However their compounds can be found used in our everyday lives. Lithium carbonates can be used in medicines, sunglass lenses and to produce metal tubes inside a television, whilst sodium carbonates can be used in tanning leather, sewage treatment and inside washing powders and liquids. Group 1 metals also have some very unusual properties.
It is unusual for metals to be soft or have low melting points, as metals, they also have very low densities. Lithium, sodium and potassium are even able to float on water! Typical metals have much higher melting points and densities.
Reaction of Alkali Metals Alkali metals are the most reactive group of metals of the Periodic Table. When reacting with water, alkali metals give off hydrogen and turn the solution from neutral or slightly acidic water to alkaline. Sodium gets hot enough to melt itself, and potassium heats up high enough to light the hydrogen gas and burn with a lilac (pale purple) flame.
Alkali metals produce 1+ ions and are more reactive going down the group.
All the reactions of an element in one group usually tell us that the rest of the group will have similar results. For example;
Alkali metals also react well with non-metals. E.g.
Group 7 – The Halogens 17-Nov All halogens form diatomic molecules, or ‘two-atom’ molecules, such as F or Cl . They are also all non-metals and contain some patterns going down the group: - They get darker in color - Melting and boiling points increases
Chlorine is used as bleach:
Chlorine + Water → Hydrochloric Acid + Chloric acid (bleach)
Cl (g) + H O (l) → HCl (aq) + HOCl (aq)
An experiment This is performed in a fume-cupboard. The chlorine gas is made by dropping hydrochloric acid on to sodium chlorate (I). The iron wool is heated strongly as the chlorine passes over. The iron and chlorine react vigorously, and red/brown iron (III) chloride is made in the reaction.
Iron + Chlorine → Iron Chloride 2Fe (s) + 3Cl (g) → 2FeCl (s)
Halogens react with metals to form salts, given that halogens are poisonous, reactions of metals and halogens should be carried out in fume cupboards. In this experiment, the chlorine has changed to chloride. We say that Halogens form halides, such as bromide, iodide and fluoride.
Unlike the alkali metals, the reactivity of the halogens decreases down the group – as it is easier for them to react with a stronger force of attraction.
The Transition Metals The transition metals lie between group 2 and 3 in the periodic table, they are very useful. The most well known as transition metals are iron/steel, copper, zinc, and gold. Transition metals are ‘typical metals’, as they are: - Hard and dense - Shiny - Good conductors of heat and electricity - Malleable and ductile - Have high melting points and boiling points (except for mercury)
Iron, cobalt and nickel are magnetic metals.
Many transition metals, such as copper, react slowly with oxygen and air. This also applies to reactions with acid and water; they react slowly, if at all. Transition metals are less reactive than group 1 metals and have a reduced activity across the periodic table.
Transition metals also form colored compounds, most group 1 compounds, however, are white. In a reaction, most transition metals form compounds which can have more than one formula. For example, there are more than one form of copper oxide. E.g. CuO and Cu O. These compounds are different colors. The roman numbers in the names can tell us which is which.
Copper (II) oxide is CuO.
Copper (I) oxide is Cu O
Catalysts speed up reactions without becoming used up. Transition metals and their compounds are important catalysts, and are highly valuable in industry.
Alloys and Corrosion 29-Nov Metals can rust; iron is one of those metals. Rust forms on the surface of iron and is a soft, crumbly substance which can flake off and cause more rusting, slowly eating away the iron/steel. Salt and acid rain also speeds up the corrosion of metals. This costs people lots of money every year. However, aluminum is a metal which resists corrosion. This is because of the tough layer of aluminum oxide on its surface which protects it. Aluminum oxide is made from reacting with oxygen and water, therefore it cannot rust.
Iron + Water + Oxygen→ Hydrated Iron (III) Oxide
Fe (s) + H O (l) + O (g) → Fe O (s) + H O (l)
The process of rusting can also be called oxidation because it is the process where a metal reacts with water and oxygen to form a metal oxide.
Preventing Corrosion/Rust Carbon is important in the extraction of iron. A giant blast furnance is used to get iron from its ore. The iron collected from the blast furnance is full of impurities. Due to this, most of it is turned to the much tougher, steel.
Ways of Prevention
There are two types of methods in preventing rust; barrier methods and sacrificial methods.
Adding other metals We can add other metals to give steel specialized properties. For example, stainless steel is made by adding chromium and nickel. Stainless steel does not rust.
Adding other metals is one way of preventing rust. It is not really a barrier or a sacrificial method.
We can coat iron or steal with: - Paint - Oil or grease - Plastic - A less reactive metal or more reactive metal
Coating is a barrier method; they help prevent the coated iron/steel from rusting.
Galvanizing Zinc is also often used to protect iron. It is more reactive than iron, therefore it will react with oxygen and moisture in the air first. Even if the coating is scratched, the iron does not rust. We say that the iron is galvanized; this is called sacrificial protection, which is a sacrificial method.
*Magnesium can also be used in place of zinc for harsher conditions.*
Electrical Protection A small electric current passing through steel will also protect it from rusting. This is possible because almost all metals conduct electricity. This is a barrier method as well.
Adding other metals to create special properties is also a way to prevent rust.
Mixtures of metals are called alloys; steel with less than 5% of other metals are low alloy steels. Those with 25% of other metals, such as stainless steel, are high alloy steels. Steel is used more than any other metal; it is valuable in the building industry. Steel is made mainly of carbon; a high amount of carbon can make the steel very hard. However, as carbon is not a metal, this steel is not an alloy.
More on: Displacement Reactions and Reactivity
Carbon is graphite. In reactivity, reactive metals want to chemically join with non-metals. Once a compound is formed, it is hard to change them back into a metal, but less reactive metals are easier to persuade as their compounds are easier to break down.
In the reactivity series above, carbon is placed below aluminum but above zinc, this means that carbon can displace any metal below aluminum.
1. What is sublimation?
2. What is the difference between diffusion and dilution?
3. i) What filtering method would I use to check the colors of dye used on an M&M? ii) How does this method work?
4. Draw any atomic structure with 6 labels.
5. Draw any electron structure.
6. What is an ion?
7. What is the AuFau principle?
8. Define: i) R.A.M ii) atomic number
9. If I had a mixture with 2 sodium isotopes where 55% is sodium-5 and 45% is sodium-15, what is the relative atomic mass of the whole sample?
10. What is the difference between a reactant and a product?
11. What is the difference between a group and period of the periodic table?
13. List 3 properties for: i) metals ii) non-metals
14. Why do group 1 elements increase in reactivity down the group but group 7 elements increase in reactivity?
15. Explain the difference between ionic and covalent bonding.
16. What is a REDOX reaction?
17. Give 3 properties for alkali metals
18. What charge do Group 1 ions usually to have?
19. What happens to the solution when alkali metals react with water?
20. Define isotopes
21. Define ions
22. What happens in a displacement reaction?
23. What are diatomic molecules?
24. What are Halides?
25. Transition metals are typical metals, list 5 properties for a typical metal.
26. What are catalysts?
27. What is corrosion?
28. What is an alloy?
29. Name two ways to prevent corrosion.
30. Why are metals able to conduct electricity?
Atom – Smallest particle which can not be broken down any further by chemical processes.
Compound – Two or more types of atoms chemically bounded together.
Element – Substance made of only one type of atom.
Mixture – Two or more chemical substances not chemically bounded together.
Molecule – Two or more atoms chemically bounded together.
Solids - Fixed arrangement - All touching - Least energy
Liquids - No fixed arrangement or pattern - All touching - Medium (energy)
Gas - Not touching - Most energy
Weakening or reducing the concentration of a substance with water or a thinner solution. The solution then turns from a higher concentrated solution to a lower concentrated one.
The exchange of liquid particles or substance from a high-concentrated area to a low concentrated area. E.g. when a perfume is sprayed into the air and the smell starts spreading across the room.
Chromatography – The process used to separate mixtures based on differences in absorbency.
Crystallization – The formation of crystals of a substance.
Distillation – The process of purifying liquid by evaporating it through boiling and recollecting it through condensing its vapors.
Filtration – The process where fluids pass through a filter.
Isotope – Atom with the same number of protons, but different number of neutrons.
Sublime – When a substance can turn from gas to solid without passing through the liquid state, and vice versa.
Relative Atomic Mass
Protons + Neutrons
Inner most shell filled first.
The Relative Atomic Mass or R.A.M of an element considers the element’s different isotopes. R.A.M considers the different proportions of each isotope in the natural mixture, for example:
If given a sample of chlorine gas and the mixture contains of 2 chlorine isotopes where: 75% is chlorine-35 and 25% is chlorine-37.
If 100 chlorine atoms: 75 = chlorine-35 and 25 = chlorine-37.
Total mass of 100 atoms: (75 x 35) + (25 x 37) = 3550
Therefore; 3550/100 = 35.5.
Relative Atomic Mass of chlorine sample: 35.5
Large enough particles of solid cannot fit/pass through the tiny holes of filter paper; therefore they remain on top of this paper.
State Symbol: what state the substance is in;
More atoms of the particle... should not change when balancing equations as it will change the chemical property and become a different chemical substance.
More of the particle; in this case, there are two particles of magnesium.
To balance an equation the number of particles on each side should be equal. E.g. there are two particles of magnesium on both sides.
3 4 5 6 7
The last group (group 8) is where the maximum number of electrons on an electron shell is reached, but if we add another shell there would be zero electrons, therefore this group is sometimes called group 0.
A mixture of water and sand is present in the test tube, the solution is poured through a funnel with filter paper and the sand is collected. The water passes through and into the second test tube below.
A solution of alcohol and water is being heated. The alcohol rises and enters the condenser surrounded by water, causing it to condense. The final product is then recollected.
The heated alcohol evaporates before water; therefore it enters the condenser first. It is then cooled by the extra water surrounding the tube. This causes condensation, which is the process where a gas changes its state into a liquid. The alcohol is recollected as a liquid.
An aqueous (solid dissolved in liquid) is present. Whilst left to settle, the liquid is evaporated whilst the solute grows into pure crystals.
Crystallization is based on the principles of solubility; a compound (solute) is more soluble in hot liquids (solvents) than cold liquids. If the solvent is cooled, the solute becomes insoluble of the solvent and forms pure, solid, crystals.
The color dye of a brown M&M (mixture) is dropped onto chromatography paper. With the paper barely touching edge of the water (solvent), the brown color gradually rises, revealing any other colors present from the original mixture.
When a solvent diffuses, it ‘drags’ whatever is dissolved in, along with it. The further the solvent diffuses, the slower the solution travels. Therefore the larger molecules fall out or halt before the smaller molecules. (Such as when smaller rocks are carried further down a river than large boulders)
The different dissolved materials separate, revealing what was once dissolved in the solution.
The effect is the same for both liquid and gas chromatography.
The Periodic Table
Amphoteric – Containing the characteristics of both an acid and a base and is capable of reacting as either substance.
Ductile – Stretchy - easily drawn into wire or hammered thin.
Electron shell – An “orbit” followed by electrons around an atom’s nucleus.
Ions – An atom or molecule where the total number of electrons is not equal to the total number of protons, giving it a net positive or negative electrical charge.
Malleable – Can be hammered or pressured into various shapes.
Metalloid – An element that behaves sometimes like a metal but sometimes like a non-metal.
Periodic Table – A tabular display of chemical elements according to the atomic number as based on the periodic law. (Mendeleev’s table)
Redox reactions – A reversible chemical reaction where oxidation and reduction occurs.
Sonorous – Can produce sounds.
Labeling electronic structures:
Na – 2, 8, 1
This means that there are two electrons on the 1st electron shield, eight on the 2nd, and one on the 3rd.
When a solid such as Iodine can change from solid to gas and does not become a liquid, it is called sublimation.
As an example, sodium is an alkali metal which is normally stored in oil. This is to keep out oxygen and moisture in the air, so to prevent the sodium from oxidising and producing sodium oxide, or rust. It is very easy to cut sodium with a knife and the inside of the cut metal is usually very shiny; as the metal may seem dull on the outside at first from minor oxidisation. Sodium melts at 98°c.
The area to the right of the borderline is where all the non-metals of the periodic table are.
The general name for non-conducting materials is called insulators. Non-metals are insulators, as they have no free electrons.
Ionic bonds form between a metal and non-metal.
Covalent bonds form between two non-metals. In covalent bonding, the elements share their electrons.
This is when the inner shells of electrons shield the outer electrons from the pull of the nucleus. The force of attraction is weakened.
All elements of the same group form ions with the same charge. This can be labelled as Li+, Na+, Cl-, F- or Mg2+.
Compounds formed by group 1 metals are usually soluble white solids. (They dissolve in water to form colorless solutions.)
Alloys – Mixture of metals.
Catalyst – Substance which accelerates a chemical reaction without letting itself be affected by the reaction.
Halides – Salt of a halogen acid.
Oxidation – The gaining of oxygen to a compound with a loss of electrons; occurs accompanied by reduction.
Reduction – The loss of oxygen to a compound with a gain of electrons; occurs accompanied by oxidation.
Salt – An ionic compound resulting from a chemical reaction, or the result of an acid and base.
A Negative Ion
Chlorine is more reactive than bromide; therefore it is capable of displacing bromide out of a solution. To do this we can simply place a chlorine solution to another solution of potassium bromide. Bromine will be displaced.
Chlorine + Potassium Bromide ’! Potassium Chloride + Bromine
Cl (aq) + 2KBr (aq) ’! 2KCl (aq) + Br (aq)
Zinc metal powder and copper oxide powder is burnt with a bunsen burner, eventually the solution b→ Potassium Chloride + Bromine
Cl (aq) + 2KBr (aq) → 2KCl (aq) + Br (aq)
Zinc metal powder and copper oxide powder is burnt with a bunsen burner, eventually the solution burns with a bright white flame and produces a white, powdery substance, zinc oxide and a brown/red powder of copper metal.
Zn (s) + CuO (s) → ZnO (s) + Cu(s)
This is an example of a displacement reaction. Due to zinc being more reactive, the metal is able to displace the copper, leaving the copper out on its own. If the reaction was the other way round, (copper powder and zinc oxide powder) no reaction would have occurred. This is because copper is less reactive.
Cu (s) + ZnO (s) → No Reaction
Displacement Reactions – Reaction where a substance displaces and sets free the element from a compound.
Reactivity – How reactive something is.
Graphite – What is used as a lubricant and moderator in nuclear reactors.
Working out the Formula
A sodium ion would have a 1+ charge, and a chlorine ion would have a 1-. 1 + (-1) = 0 therefore only one atom of each compound is needed. The formula is HaCl.
If the sodium reacted with oxygen it would make sodium oxide. Sodium has 1+ charge and oxygen has 2- charge. 1 + (-2) = -1! Therefore we need another sodium ion to make the charge neutral. 2 + (-2) = 0. The formula is Na O.
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