It’s everything chem ppt !!

1. It’s everythingchemppt!! Mr Fedley

2. Let’s Explode Some Stuff • Yes – we are finally exploding some stuff. • Elements in the first Period (Row) of the Periodic table react violently with water. • Today, we will look at Sodium (Na) and Potassium (K). • Please be careful and listen to my instructions. This is dangerous!!

3. Basic Chemistry Ideas http://www.learner.org/courses/essential/physicalsci/session2/closer3.html • Everything is made out of small particles that are too small to be seen. • These small particles always move around. How much they move around depends on if they will be a solid , liquid or gas. http://blackcoolhairstyle.blogspot.com/2010/06/mercury-liquid.html http://inyoprocess.com/white-papers-mainmenu-5/9.html?task=view

4. Solids, Liquids and Gases • If a particle is heated, then it becomes more excited and moves around more. • When a particle moves around more, their bonding (how they are held together) can become broken and they change state (eg Liquid to gas) • The opposite happens if a particle is cooled. That is, the particles are less excited and move less. They can form new bonds and they change state (eg gas to liquid)

6. What Are Atoms? • According to the dictionary, atoms are ‘the smallest particle of an element that can exist either alone or in combination’. • In English, an atom (原子) is the smallest thing that we can see under a microscope. Everything is also made of atoms.

7. Some Deciphering • Everything is made up of particles (粒子). These particles are made up of different atoms that have different properties (プロパティ). • You can almost think of atoms as ingredients in food. We have 100 or so ingredients. • When we make a cake, it will have different properties depending on it’s ingredients. http://www.frequentbakes.com/our_menu.htm http://goodstufffood.com.au/products.php?product=SWE&subproduct=CAKES

8. The Periodic Table – Our Ingredients

9. Differences In Atoms • There are over 100 atoms. • Every atom is different as it contain a unique number of neutrons (中性子), protons (陽子) and electrons (電子). • An electron is found in an electron shell and has a negative charge (負電荷). • A proton is found in the nucleus (核) and has a positive charge (正の電荷) • A neutron is also found in the nucleus and has no charge (ない手数料なし).

10. Atoms Atoms are the “building blocks” of all matter and are the simplest form of molecule. They are all made up of protons, neutrons and electrons. They look something like this: The Nucleus – this contains neutrons and protons Electrons – these orbit around the nucleus

11. Atomic & Mass Number

12. Electron Shells • In a stable atom, the atomic number tells us how many electrons or in the atom. • Atoms are found in electron shells (電子殻) which take up most of the space of each atom. • Different electron shells can hold different amounts of electrons • Shell 1 holds up to 2 electrons • Shell 2 holds up to 8 electrons • Shell 3 holds up to 8 electrons

13. Electron Shells Continued • The shells in in order, starting with the inner one (lowest energy level) • All the elements in a group have the same number of electrons in the outer shell • The group number is the same as the number of outer electrons • Period 0 (Noble Gases) have full shells. • The period number states how many shells there are.

14. Electron Dot Diagrams • Each atom normally has the following: • A nucleus which contains • Protons which have a positive charge. This is the same as the atomic number • Neutrons which have no charge. This can be worked out by subtracting the molecular number from the atomic number. • Electron shells which contains • Electrons which have a negative charge. This is the same as the atomic number

15. Electron Shells • Electrons go into electron shells. They like to be as close to the nucleus as possible. As such, they fill up the closest shells first. • The following shells can hold: • 1st shell: up to 2 electrons • 2nd shell: up to 8 electrons • 3rd shell: up to 18 electrons unless it’s the last shell. Then it’s 8 • 4th shell: up to 32 electrons unless it’s the last shell. Then it’s 8.

16. Catatonic Compounds • So what’s the difference between elements and compounds? • This is what we will now be learning about. We’ll be seeing how different atoms bond and what properties then might have because of that. • For example, does anyone know why most non-metals when bonded have a low MP and BP.

17. Elements & Compounds • Elements are one type of atom by themselves which is pure. • A compound is two or more atoms together. • Noble Gases normally don’t form compounds • Atoms like to have full outer electron shells. • If there is 1, 2 or 3 electrons in outer shell, it is more likely to lose those electrons • If there is 5, 6 or 7 electrons in outer shell, it’s more likely to gain electrons.

18. Ions • Ions are atoms that have gained or lost electrons to have a full outer shell • Sodium loses one electron so it now has a positive charge. So we write it as Na+ (the plus is normally in the same place as a power in maths). • So, if you lose 3 electrons it will be 3+. • If you gain 3 electrons it will be 3-.

20. Some Hints • When you write the charge (+ or -) of an ion, you write it in the same place as a power. Eg Na+. • When you write how many atoms you have, you write it down on the bottom right. Eg O2. • When making a compound, it is normally stable. This means there is the same number of electrons as there are protons.

21. Ions • Atoms like to have full outer electron shells. As such, they tend to gain or lose electrons. • Atoms with 1,2 or 3 electrons in the outer shell (apart from Hydrogen) tend to lose electrons and become a positive ion (positive valency) as they have more protons than electrons. • Atoms with 5, 6 or 7 electrons in the outer shell tend to gain electrons and become a negative ion (negative valency) as they have less protons than neutrons.

22. Ionic Bonds • When an anion (negative ion) and cation (positive ion) are close, they can bond or join together. • For example in Sodium (Na+) and Chlorine (Cl-), the Sodium will transfer one electron to Chlorine. • Both atoms are now stable and have bonded together.

23. A Special Bond • The total charge in an ionic bond is normally 0 • 2 Chlorines (Cl-) can bond with one Magnesium (Mg 2+). • An ionic compound occurs when an anion and cation bond together. • When they bond together – think of them as little magnets. + and – want to go next to one another as 2 of the same charges repel.

25. Naming anions in Ionic Equations • By itself, Chlorine is called Chlorine. • As a compound, it’s called Chloride. • Anions (atoms with more electrons than protons) change their name with ide on the end. • Examples are: • Oxygen = Oxide • Iodine - Iodide

26. Cations • An Ion is an atom that has gained or lost electrons to have a full outer shell. • Elements from Group 1, 2 & 3 normally lose electrons and will have more protons than electrons. As such, they have a positive charge. They are called cations. • Ions in Group 1 are charged +1 • Ions in Group 2 are charged +2 • Ions in Group 3 are charge +3

27. Anions • When a metal and a non-metal bond together, it’s called an ionic bond. They gain or donate electrons – they do not share them. • Elements from Group 5, 6 & 7 normally gain electrons and will have more electrons than protons. As such, they have a negative charge. They are called anions. • Ions in Group 7 are charged -1 • Ions in Group 6 are charged -2 • Ions in Group 5 are charged -3

28. Speeding Up Reactions • Speeding up reactions are quite important. If we don’t, then it may take too long for the reaction to occur. Some ways to increase the speed of a reaction includes: • Change temperature (normally increasing it) • Change the concentration of the reactant (ingredient not product) • Increase surface area • Helper Chemicals (catalyst)

29. Change Temperature • Today, we are going to start looking at how temperature can change reaction rates. • When you increase the temperature, it can speed up physical and chemical reactions. • This is due to particles moving more quickly due to a higher temperature. • If the particles move more quickly, they are more likely to collide and speed up the reaction.

30. Temperature Change • When you increase the temperature, it can speed up physical and chemical reactions. • This is due to particles moving more quickly due to a higher temperature. • If the particles move more quickly, they are more likely to collide and speed up the reaction. • An increase in temperature can also allow more solid to be dissolved into a liquid. This is a physical change. • Please watch this space as there will be more experiments for temperature change but we need bunsen burners 

31. The Mole • The mole is used to tell us the mass of a particular number of atoms. • There are 6.02 X 10^23 atoms in a mole. • When you look at the periodic table, the molecular mass tells you how much 1 mol of that atom weighs. • For example, 1 mol of Hydrogen = 2 grams I was not talking about this mole!! http://crashofcourse.blogspot.com/2010/06/mole-attack.html

32. The Mole in Liquids • In Chemistry, you might notice that liquids have 1M, 10M, 0.1M on them. • This tells us how many moles of an element or compound are dissolved in 1L of water. • As such, 1M is 10 times stronger than 0.1M etc. http://www.thinkgeek.com/tshirts-apparel/unisex/sciencemath/e635/

33. Concentration • Concentration refers to the amount of dissolved solid in a solution. • The more concentrated a solution, the more of that particular element of compound has been dissolved. http://diylol.com/meme-generator/chemistry-cat/memes/do-you-know-any-jokes-about-silcon-si

34. Speeding Up A Reaction • During a chemical reaction, the number of particles is important. More particles in the same area = faster reactions most of the time. • As such, a 10M Hydrochloric Acid solution should be more reactive than a 0.1M solution. • When you dissolve more solid into the same amount of liquid, there is a greater chance of a reaction. • For example, heating water to dissolve more sugar could be very beneficial if we have a reaction that requires dissolved sugar.

35. Chemical Reactions • There are many different types of chemical reactions. These Include • Oxidation • Combustion • Reduction • Decomposition • Neutralization (acid-base) • Precipitation • Displacement

36. Oxidation Reaction • Oxidation occurs when a substance reacts with oxygen: • Pb (s) + O (g) PbO (s) • Some examples of oxidation includes rust and when food de-colors such as apples and butter. • Chem hint: in an equation (s) = solid, (l) = liquid, (g) = gas & (aq) = aqueous or in water

37. Combustion • A combustion reaction occurs when a substance burns in oxygen. • Some good examples include burning fossil fuels or wood. • Methane + Oxygen  Carbon Dioxide + Water • Hydrogen + Oxygen  Water

38. Reduction Reaction • A reduction reaction occurs when a substance loses oxygen. This normally requires a large amount of energy. • Iron Oxide + Carbon  Iron + Carbon Dioxide • Reduction reactions are commonly used in the preparation of ores into a metal.

39. Decomposition Reactions • Decomposition occurs when a compound breaks down into simpler substances. • Not all compounds break down easily, but some can and this can be brought about through heat, light or electricity. • Silver Oxide  Silver + Oxygen

40. Neutralization • Neutralization is the reaction between an acid and a base (alkali). They normally react to form a salt and water. • Hydrochloric Acid + Calcium Carbonate  Calcium Chloride + Water + Carbon Dioxide • We will be learning more about this topic after this unit.

41. Precipitation • Precipitation occurs when two aqueous solutions are put together and a solid forms. • The precipitate can often be strange colors – but many are just white (yay!) • Lead Nitrate (aq) + Potassium Iodide (aq)  Lead Iodide (s) + Potassium Nitrate (aq)

42. Displacement • A displacement reaction occurs when a more reactive element takes the place of a less reactive one in a compound. • Chromium Oxide + Aluminium Chromium + Aluminium Oxide

43. Different Types of Reactions • Exothermic Reactions – These types of reactions release energy. Energy might be released in the form of heat, light or sound. They often give off heat and may feel hot. • Endothermic Reactions – These type of reactions require energy. Energy may be absorbed in the form of heat, light or sound. They normally absorb heat and may feel cold.

44. Combustion Reactions • Combustion reactions need oxygen to occur. They often have water as one of the products. • Remember reactants are on the left of an equation, the products are on the right of an equation. • They are normally exothermic and release energy in the form of heat, light and sound. • 2 C8H18 + 25 O2 18 CO2 + 18 H2O

45. Stating States • When we write equations, we normally state which type of state (solid, liquid, gas) it is. They are stated below: • (l) = liquid • (g) = gas • (s) = solid • (aq) = aqueous (in solution or dissolved)

46. Oxidation Reactions • Oxidation occurs when a substance reacts with oxygen. An oxygen atom is gained in the molecule or compound in question • 2Mg (s) + O2 (g)  2 MgO (s) • Note how the Magnesium now has one more oxygen atom on the right hand side. • The Magnesium has now been oxidized

47. http://scienceprojectideasforkids.com/2010/oxidation-sparklers/http://scienceprojectideasforkids.com/2010/oxidation-sparklers/

48. Reduction • A reduction reaction occurs when a compound or molecule loses oxygen. • CO2 (g) + C (s)  2CO (g) • In this reaction, note how the Carbon Dioxide loses one oxygen to become Carbon Monoxide. • CO can be used to kill someone quite easily as your lungs thinks it’s oxygen and you choke to death 

49. Displacement • When two metals can react, the metal that is most reactive will react first. • This quite often happens with Zinc bolts on ships. Zinc is more reactive than Iron, so the Zinc with oxidize first before the Iron. • Iron Oxide + Carbon  Iron + Carbon Dioxide • They are also often used in find halides (Group 7) in solutions as they will be displaced in Silver Nitrate. • We will learn a bit more on this next lesson

50. http://www.frankswebspace.org.uk/ScienceAndMaths/chemistry/reactivitySeries.htmhttp://www.frankswebspace.org.uk/ScienceAndMaths/chemistry/reactivitySeries.htm