revision: Chemistry - Topic 2

I will update slowly. I'm sorry. Like Biology, if there's a topic you want me to cover ASAP, comment below or fb inbox me or whatever is good for you s'good for me :D
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Stay tune! I'll try to update as soon as possible but I'm quite busy now...

I've been getting a lot of people saying they want to contribute...if you do...again, comment, inbox me whatever and I'll add them to my notes.... Apparently it helps others...so thankyou. For being my motivation.

So...here's my revision...Yes, I'm beginning early thanks to the people who requested.

UPDATES = NEW NOTES.

There are 3 main types of chemical bonding:
IONIC : transfer of electrons.
COVALENT : sharing of electrons.
METALLIC: when a metal bongs with a metal.
"A lattice of positive ions is electrostatically attracted to a "sea of delocalized electrons."

IONIC BONDING.
-Occurs when metal bond with non metal.
-Transfer of electrons from on to the other resulting in a positive or negative charge.
-Cation = positive ions, anion = negative ions.
-An electrostatic attraction between oppositely charged ions.

Metals in group 1,2,3 can get a full outer shell by losing electrons becoming a positive ion. A cation. Remember -> Ca+ion
This process is called oxidation (OIL RIG = Oxidation is loss, reduction is gain.)

Non-metals in group 5,6,7 can get a full valence shell by gaining enough electrons resulting in a negative ion. An anion. This process is called reduction.

Once the ions are formed, they are oppositely changed and will attract one another. This forms a GIANT IONIC LATTICE.

Giant ionic lattices have a lot of strong ionic bonds and need a lot of energy to be broken...
-This causes them to have high  melting points. The stronger their bonds, the higher their melting points.
-The do not conduct electricity unless in solution (aqueous) or when molten, since then, the ions are free to move and can carry charge.
-They are usually soluble in water (as water is a polar molecule = one end is negative and one end is positive and they can cluster around the ions allowing them to separate, overcoming the strong attracting forces that hold the lattices together.

COVALENT BONDING.
When two non-metals combine, they form a covalent bond. As they both need to gain electrons, they do this by sharing electrons.

-Attraction of positively charged protons in the nucleus and the negatively charged electrons on the shells that are shared between two.
*YOU GOTTA BE ABLE TO DRAW DOT AND CROSS DIAGRAMS, BRO!!!*

SMALL MOLECULAR STRUCTURES-
-They have small intermolecular forces, and require little energy to break these attractions therefore they have low melting and boiling points.
-They don't conduct electricity.
-Dissolves  in non-polar solvents.

GIANT COVALENT STRUCTURES-
Non-metals can form up to three or four bonds, and it is possible to form giant structures linked by covalent bonds.
-For example, there are two forms of carbon.
Diamond is the strongest natural substance, which is covalently bonded forming a crystal.
Graphite consists of hexagonal layers which can slide above each other - so we can write with graphite :D :D :D
-They both require a lot of energy to break bonds as it is necessary that all are broken for these two substances to melt - therefore, they have extremely HIGH melting points.
When elements are found to exist in more than one crystalline form they are known as ALLOTROPES.
-A third allotrope of carbon is Buckminster which has a football/soccer ball shaped structure.
PROPERTIES:
-Hard,
-High melting points.
-Do not conduct electricity because they don't have charged particles.
-Insoluble.

METALLIC BONDING-
Metallic bonding means giving up their outer electrons and sharing :) BECAUSE SHARING IS CARING. So they share with their neighbors. This creates a bunch of delocalized electrons and we end up with a SEA OF DELOCALIZED ELECTRONS! The positive particles (ions) then pack themselves as tightly as possible forming giant metallic structures.
-They conduct electricity because the delocalized electrons are free to move.
-They have high melting points as the structures have strong attractions and require a lot of energy.
-Malleable and ductile because the ions can slide over each other.

2.5 - understand that the noble gases (Group 0) are a family of inert gases and explain their lack of reactivity in terms of their electronic configuration.

"The Noble gases were called 'inert gases' because they are nonreactive. The lack of reactivity is explained by the full outer shell called the 'duplet' in the case of helium and 'octet' in the case of the others. The full outer shell gives stability to the noble gas atoms so they do not need to react. This group an be called group 8 in view of 8 electrons in the valence shell, but helium is an exception having a duplet."

-The elements of group 8/0 are the noble gases (He, Ne, Ar, Kr, Xe, Rn)
-They are inert and almost completely unreactive.
-This is because they all have 8 electrons on their outer valence shell.
-When chemical reactions occur, they form compounds to have full valence shells - becoming more stable.
-This can occur by ionic bonding (transferring electrons) or covalent bonding (sharing electrons).

1.28 - describe the formation of ions by the gain or loss of electrons.

"Students should be able to explain the ion charge from the imbalance in the number of electrons and protons as compared to the atom. The number of electrons lost or gained depends on the number needed to gain a full outer shell and a stable 'noble gas' configuration. The terms 'cation' and 'anion' should be known."


Ions are atoms or molecules with an electric charge due to the gain or loss of electrons (Yes, they do this to be like noble gases - they want to have the same, amount of electrons, like we want to have the same amount of clothes or shoes as those celebrities :D) REMEMBER -> These form ionic compounds, the transfer of electrons between metals and non-metals.

·      The loss of electrons from an atom forms a positively charged ion (cation) – oxidation

If electrons are lost, the ion then has a positive charge and depending on the amount of electrons lost, they are written with a + sign.
So...if an ion A loses 2 electrons it is written as A^2+
These are knowing as cations. Remember -> Cation, CA+ion. See? it's a + sign.
The elements from group 1 -3 will most likely form cations as they will lose electrons.

·      The gain of electrons by an atom forms a negatively charged ion (anion) – reduction

Vice versa, if electrons are gained, the ion then has a negative charge and these are written with a - sign.
If ion A gains 3 electrons the it is written as A^3- because it has a negative charge of 3!
And they are known as anions. Remember -> Anion or A-negative-ion = Anion, negatively charged ions.


Elements from group 5-6 will form anions as they will gain electrons. Group 8/0 are noble gases and are inert so they do not form ions. (They are too cool. Celebrities -.-)

These are known as ionic compounds! The transfer of electrons due to a reaction between metals and non-metals.

1.29 - understand oxidation as the loss of electrons and reduction as the gain of electrons.

Remember -> OIL RIG is Oxidation Is Loss, Reduction Is Gain.

1.30 - recall the charges of common ions in this specification.
"The charges on the simple ions can be worked our from the group number and hence the number of valence electrons in the case of elements later in groups 1-3 and 5-7. Group 4 elements are largely covalent (see later) except that lead and tin can from +2 and +4 ions (see objective 1.21)"

·      Ionic bonding between metals and non-metals forms very strong bonds of electrostatic attraction between these positive and negative cations and anions

The number of ions combined in order to cancel out, balance or neutralise the charges determines the formula of the compound 

Metallic Cations:                   Non-metallic Anions:                                         Polyatomic Ions:

Gp 1    Gp 2    Gp 3                Gp 5                Gp 6                Gp 7                            Learn These:

Li⁺       Be²⁺     B³⁺                   N^3- (nitride)     O^2- (oxide)       F^- (fluoride)                 NH₄⁺ (ammonium)

Na⁺      Mg²⁺    Al³⁺                                         S^2- (sulphide)   Cl^-(chloride)                OH^- (hydroxide)

K⁺        Ca²⁺     Ga³⁺                                                                 Br^- (bromide)               NO₃^- (nitrate)

Rb⁺      Sr²⁺      In³⁺              (Note; All ‘...ide’ compounds)     I^- (iodide)                    CO₃^- (carbonate)

Cs⁺      Ba²⁺     Tl³⁺                                                                                                      SO₄^- (sulphate)

Fr⁺       Ra²⁺                                                                                                                 PO₄^3- (phosphate)

Transition Metal Cations: (Charge is shown by Roman numeral or Need Learning)

Zn²⁺ / Ag⁺ / Copper I: Cu⁺ / Copper II: Cu²⁺ / Iron II: Fe²⁺ / Fe III: Fe³⁺ / Nickel II: Ni²⁺ etc.

1.31 - deduce the charge of an ion from the electronic configuration of the atom from which the ion is formed.

1.32 - explain using dot and cross diagrams the formation of ionic compounds by electron transfer, for combinations of elements from Groups 1, 2, 3 and 5,6,7.

BY RACHEL

Hey y’all, this is Rachel. :) I don’t really have the pen tools and stuff for diagrams like Malisa’s, but I’ll try to simplify everything as best I can with words.

1.35 understand the relationship between ionic charge and the melting point and boiling point of an ionic compound

Ionic structures have really, really high melting points. You could heat them over a Bunsen burner for ages and not get anywhere. This is because there are millions of electrostatic attractions that need to be broken. Ionic structures are giant structures.

 ß this is an ionic structure.

Okay, getting to the point. This ‘objective’ is to understand why ionic charge affects the melting point and boiling points of ionic compounds. We know that these are high to start with anyways, but some compounds have even higher melting points than others.

Important: if you’re going to remember anything from this objective overview, this next thing would be it, because it’s most commonly brought up in tests and exams. You are likely to be asked about magnesium oxide and sodium chloride, and told to explain which one has a higher melting point and why.

Magnesium oxide is MgO, and sodium chloride is NaCl.

Magnesium is in group 2, so it loses two electrons, giving it a positive charge of 2. (Mg^2+). Oxygen is in group 6, so it gains two electrons, giving it a negative charge of 2. (O^2-) [The ^ means power, by the way, In simple English, it’s the small numbers above the symbol.]

Sodium (Na) is in group 1, so it only loses 1 electron, meaning it only has a positive charge of 1. Chlorine (Cl) is in, you guessed it, group 7, so it only gains 1. No one writes 1+ next to it because it seems redundant, but for this question, you need the charges for MgO and NaCl engraved into your heads.

Anyways,

 Magnesium oxide (MgO) has the higher melting point, because it has an attraction that is twice as strong as that of sodium chloride (NaCl), making the electrostatic attractions of MgO harder to break.

Think of it like….bigger charge, bigger attraction.

1.36 describe an ionic crystal as a giant three-dimensional structure held together by attraction between oppositely charged ions

I’ve already shown you that diagram of a giant lattice, just scroll up to the green and purple balls that form a square. To describe this, you just need to know that each positive ion (cation) has 6 negative ions (anions) around it, the same way each negative ion has 6 positive ions around it.

Think of that really lame, cheesy excuse saying that people say – opposites attract. That may not be a legitimate rule in real life, but it is in science. Think of magnets – the blue end of one magnet is attracted to the red end of another. Opposites attract.

Now imagine you’re building a ball and stick model of this, and your teacher hasn’t told you how big to make it. You’re really bored, and you have all of class, so you keep adding, and adding, and adding…and you never really have to stop until you run out of balls and sticks. Ionic crystals are huge. They just. Keep. Going. Because ions can keep being added on. This is something else you might need to know.

1.48 understand that an electric current is a flow of elections of ions

Ionic structures conduct electricity, but only when dissolved in water (solution). Think – electrons, electricity, electrons, electricity…you don’t get electricity if there are no moving electrons.

Ionic compounds have to be dissolved to conduct electricity, because they don’t contain any mobile electrons when they are solid. When they melt or dissolve, the electrons become free to move around. This could get a lot more complicated, but this is all you need to know for the test, so I’ll end it this one here to keep from confusing you.

It might be worth it to mention that COVALENT STRUCTURES DO NOT CONDUCT ELECTRICITY. Do. Not. Forget.

1.50 understand why ionic compounds conduct electricity only when molten or in solution

I just covered this in 1.48….see? It’s not so bad. Everything overlaps. :)

1.38 understand the formation of a covalent bond by the sharing of a pair of electrons between two atoms

Shaaaaring is caaaring! ;) (It’s Rachel again obvs)

It’s like when your friend forgot to bring their text book to class (pretend we still use textbooks, okay?) They don’t have enough class equipment. You have your textbook, but you didn’t bring your laptop, so you can’t write anything down. You don’t have enough class equipment either. They did bring their laptop, and since you sit next to each other, you decide to share. You put the textbook and laptop in between the two of you, and take turns using it. Your friend now has a textbook, and you have access to a laptop. This is covalent bonding.

It happens between two non-metals. Neither are willing to give up their electrons, as that wouldn’t leave them with a full outer shell. Instead they share the electrons, so that they can both have full outer shells without finding some other element to take them from.

A simple example is water (H₂O). It is important to note that hydrogen only has one shell, so it’s outermost shell can only hold two electrons. Oxygen has 6 valence electrons, and it needs two. Hydrogen can’t really give up its one electron, or it will just be a nucleus…with no electrons.  So since hydrogen only has one electron (that, I must reinforce, it won’t give up), and oxygen needs two electrons, two hydrogen atoms scoot closer to oxygen…and closer…until their shells overlap, and the electrons can pass back and forth between the oxygen and hydrogen atom. This is often shown in dot and cross diagrams (that you need to know how to draw).

The blue X represents the electrons that originally belonged to each hydrogen atom, and the red dots represent the electrons that originally only belonged to the oxygen atom. Since they overlap now, they can share. :)

Keep in mind that H₂O is probably one of the simplest of covalent bonds, and that if you are asked to draw dot and cross diagrams for another compound, you will have to use your logic and maths skills. Use a pencil, sketch some circles out, play around with it until each atom has a full outer shell.

Another thing to note: a covalent bond in written terms can be represented by a hyphen (-) between the two atoms. For example, hydrogen fluoride can be shows as H-F.

1.40 explain, using dot and cross diagrams, the formation of covalent compounds by the electron sharing for the following: hydrogen, chlorine, hydrogen chloride, water, methane, ammonia, oxygen, nitrogen, carbon dioxide, ethane, ethene

I guess all I can do for this part is give you tips.

I’d suggest that when you first draw each atoms valence shell, draw the first four electrons in the north-south-east-west positions by themselves first. THEN add on the rest as pairs…does this make sense? For example, nitrogen has 5 valence electrons. I would start off by drawing one at 12 o’clock, 3 o’clock, 6 o’clock and 9 o’clock, so I’ve drawn 4 out of 5 electrons. I’d put the last one back at 12 o’clock next to the first one I drew, so that it’s neater. Examiners also tend to prefer this.

Also, ONLY DRAW THE OUTERMOST SHELL.  If you ever have to show a covalent bond involving an atom like…I don’t know, bromine, you’ll have a lot of shells to draw, and since our hands don’t really work as perfectly as compasses, it will be messy and really hard to read. Draw. Only. The outermost. Shell.

I’m prettttttty sure we all get periodic tables, so we can work out the diagrams on our own. At this point, they should give us the equations for these molecules, so we’re all good :)