Ammonia

10 05 2012

Ammonia NH3 contains 3 electron pair bonds (N-H) and one lone pair, represented on the model I made below as the empty tip of the pyramid shape. The blue centre represents Nitrogen, and the 3 white bits are Hydrogen.

The pyramidal shape is a result of the electron pairs (bonded and lone) repelling each other. Remember where the most repulsion occurs:

Most electron repulsion between:       lone pairlone pair

Strong repulsion between:                      lone pairbonded pair

Least repulsion between:                         bonded pairbonded pair

These repulsion rules explain why the bonded pairs are pushed downwards to form the base of the pyramid, they don’t want to be anywhere near the lone pair at the tip of the pyramid. But they also don’t want to be near each other, so they go as far as they can away without breaking their bonds to the nitrogen centre.

credit: wikimedia.org

What shape do you think the Ammonium ion NH4+ would have?

Ammonium is isoelectronic CH4 (Methane) so has the same shape, tetrahedral. Remember ‘iso’ in Greek means ‘equal’, so in chemistry it means ‘the same’ or ‘no change’, remember we studied Isotopes of elements in Lesson 2.

Methane (credit: wikimedia.org)

Here’s an ammonium ion I drew below:

NH3 + H+  = NH4+

Dative bond: Both electrons of a shared pair are donated by one of the bonded atoms.

In the example of ammonium this symmetrical ion has a dative bond that is hard to distinguish from its other covalent (electron sharing) bonds because of its symmetry. But as you can see there are two dots instead of dot cross, and the + sign atop the brackets surrounding the ion represents the electron the hydrogen atoms gives away. Electrons are negative so giving them away means you become more positive.

Ammonium is a positive ion so is called a cation. Here’s a model of amonium I made:

 

Methane:

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Lesson 4: Formulae and Equations

8 05 2012

Click to download Lesson 4: Formulae and Equations (Word document) Formatting looks much better when you download it! When pasting from Word to this post some of my tables and colours were altered.

Empirical Formula is the simplest whole number ratio of each type of atom in a compound.

Example:

Chemical Molecular Formula Empirical Formula
Glucose C6H12O16 CH2O
Ethanoic Acid CH3CO2H CH2O

Looking at the molecular formula for Ethanoic Acid you can see there are

2 Carbon, 2 Oxygen and 4 Hydrogen atoms, this whole molecular formula has then been divided by a common multiple, in this case 2, so in the Empirical Formula you have:

1 Carbon, 1 Oxygen and 2 Hydrogen atoms.

In some questions you will be asked to calculate the Empirical Formula of a compound.

As an example, let’s work out H2SO4 by breaking it into its component parts (ratio H : S : O). If the question says we have 2.45g of H2SO4 containing 0.05g Hydrogen, 0.8g Sulfur and 1.6g Oxygen. Create yourself a little table, and work out the details from top to bottom:

In basic form, your working table could look like this:

H S O
Mass (g) 0.05 0.8 1.6
Molar mass (gmol-1) 1 32 16
Moles 0.05 0.025 0.1

Then divide all moles values by the smallest; you’re left with the result of the simplest ratio of how many of each atom makes up the compound.

Let’s try another Question…

Analysis shows that 0.6075g of Magnesium combines with 3.995g of Bromine to form a compound. Find the Empirical Formula of this compound.

Mg Br
Mass (g)÷Molar mass (gmol-1) 0.6075g ÷ 24gmol-1 3.995 ÷ 80gmol-1
Moles 0.025 ≈ 0.3mol 0.5mol
Divide mol values by the smallest 0.3 ÷ 0.3 0.5 ÷ 0.3
Simplest Ratio 1 1.6666… ≈ 2
Empirical Formula Mg Br2

The simplest ratio of Magnesium is 1 so just write Mg, which is one atom. For Bromine the ratio was 2, so you have to write Br2 two atoms being the simplest ratio. Simply it is a 1:2 ratio.

Molecular Formula shows the exact number of atoms in a molecule.

It is used for compounds that form covalent molecules (covalence is where elements share electrons in a chemical bond.) Remember that Noble gases are monoatomic (mono means one, single) and other elemental gases are diatomic (di means two or double)

Diatomic example: in molecular form Hydrogen gas is H2, Chlorine gas is Cl2, and Nitrogen gas is N2.

CH4 Methane is made up of 1 Carbon atom and 4 Hydrogen atoms.

We can use the ‘octet’ rule with ionic formula (which is explained in the next heading) to predict molecular formula, but only for s and p block elements not d block. (In Lesson 1 we discussed that the noble gases all have a ‘stable octet’ of 8 electrons in their outer shell making them very unreactive, because they don’t want to give away or accept electrons).

So if we wanted to predict the molecular formula of chlorine oxide…

Cl: 8 – 7 = 1 so Cl1-

O:  8 – 6 = 2 so O2+

The charges of each element must always balance, but currently we have 1- and 2+

You cannot subtract from the 2+ or add to the 1-, the charge does not change, but the amount of an element can change. In this case we need 2 atoms of Chlorine to balance out the 2+ charge of Oxygen. So we can predict the molecular formula to be:

Cl2O

So we have Cl 1- (x2) with O 2+, balanced together so that the overall charge always equals zero. This compound is called dichlorine monoxide.

Ionic Formula

This is the formula of an ionic compound whereby the charges of the elements are made to balance.

 Notice how in each equation Chlorine has a different charge, the ionic charge of the element itself has not changed (it is still 1- in each), but each time it has needed more atoms of itself to balance with the positive charge of the element it is bonding with. So sometimes it is just Cl, other times 2 atoms of Cl are needed, Cl2, and sometimes 3, Cl3.

We will go into more details on ionic compounds, and charges on the periodic table soon.

Structural Formula

The structural formula shows the arrangement of atoms and groups in a molecule.

Essential for organic compounds.

Example

Ethanol

oxidised to

Ethanal (Acetaldehyde)

CH3CH2OH

=>

CH3CHO

For organic reactions we often write unbalanced equations showing only the structural formula of the principle organic reactant and products. (reactant Þ products).

For inorganic reaction we usually write balanced equations showing all the reactants and products.

Examples of an inorganic equation:

1) Ordinary equation

Mg + H2SO4  => MgSO4 + H2

2) Ionic equation

Mg + 2H+ => Mg2+ + H2

Hydrogen gains electrons so is reduced. (Oxidation is loss of electrons, Reduction is gain of electrons. Remember OIL RIG”: Oxidation Is Loss, Reduction Is Gain.)

Half Equations

Half equations are used to show the loss and gain of electrons in a reaction, in other words to demonstrate the oxidation and reduction.

Oxidation Is Loss

Mg => Mg2+ + 2e

The 2e are the two electrons lost by Magnesium, electrons are negative so a loss of them makes the atom more positive, and hence Mg becomes Mg2+

Reduction is Gain

2H+ + 2e => H2

In this half equation Hydrogen is initially positive, until 2 electrons (2e) are added and it becomes negative. Remember, it must balance.





Lesson 3: Mass and Moles

7 05 2012

 Download lesson as Word document, click: Lesson 3 

Mass and Moles

Objective: To explore different types of mass of elements, and be introduced to the measurement of moles.

Relative Isotopic Mass

This is the mass of an atom of an Isotope compared with 1/12 of the mass of an atom of Carbon-12.

Example: Oxygen-16 has a relative Isotopic mass of 16.0, Sodium-23 has a relative Isotopic mass of 23.0.

 

Relative Atomic Mass (Ar)

The relative atomic mass is the weighted mean mass of an atom of an element, this is calculated using the different masses and relative abundances of all the Isotopes of a particular element.

For example:

35                        37

     Cl                         Cl

17                        17

If I am told that 75% of Chlorine atoms have an atomic mass of 35, and 25% have an atomic mass of 37, I can calculate the relative atomic mass.

(35×75) + (37×25) ÷ 100 = 35.5

35.5 is therefore the relative atomic mass of Chlorine, an average taken using the types and amounts of other Isotopes of the element.

Another example:

 

24                       25                        26

    Mg                               Mg                        Mg

12                       12                          12

 

78.6%                10.1%                   11.3%       < abundance of each Isotope of the element magnesium.

(24 x 78.6) + (25 x 10.1) + (26 x 11.3)

________________________________    = 24.3

100

As you can see Mg 24 is the most common / abundant Isotope, so the final average is nearest to 24. Relative atomic mass allows for the most accurate average mass of a particular element.

What is the difference between Relative Isotopic Mass and Relative Atomic Mass?

This is simple; Relative Atomic Mass takes into account ALL the Isotopes of a particular element (as demonstrated above), whereas Relative Isotopic Mass means the mass of just ONE Isotope of a particular element.

Example: Chlorine-35 has relative ISOTOPIC mass of 35. Chlorine-37 has a relative isotopic mass of 37. But if we want the relative ATOMIC mass, we must add the isotopes, multiply by abundance, and divide by one hundred to find an average mass (this exact question was covered above).

Relative Formula Mass

Also known as relative molecular mass (Mr). Relative Formula Mass is the weighted mean mass of a molecule (compared with 1/12 of the mass of an atom of Carbon-12).

So, where Relative Atomic Mass dealt with an atom of a whole element, Relative Formula Mass deals with the mass of a molecule (a molecule is made up of two or more chemically bonded atoms).

Many elements and compounds are made up of simple molecules like N2, O2 or CO2

Compounds with giant structures do not exist as molecules, for example: Ionic compound NaCl, or covalent compound SiO2, so ‘relative formula mass’ is seen as more accurate than saying molecular mass.

Remember: Atoms, Molecules, and Compounds;

A molecule is formed when two or more atoms join together chemically. A compound is a molecule that contains at least two different elements. All compounds are molecules but not all molecules are compounds.

Molecular hydrogen (H2), molecular oxygen (O2) and molecular nitrogen (N2) are not compounds because each is composed of a single element. Water (H2O), carbon dioxide (CO2) and methane (CH4) are compounds because each is made from more than one element. The smallest bit of each of these substances would be referred to as a molecule. For example, a single molecule of molecular hydrogen is made from two atoms of hydrogen while a single molecule of water is made from two atoms of hydrogen and one atom of oxygen.” http://education.jlab.org/qa/compound.html

To calculate the relative formula mass (Mr) of a substance all you have to do is add up the relative atomic masses of all the elements.

H2O     Water there fore has a relative formula mass of 18.

Hydrogen has an atomic mass of 1, there are 2 atoms of Hydrogen in water so multiply by 2.

Oxygen has an atomic mass of 16, so 16 + 2 = 18.

What is the relative formula mass of K2CO3 ?

 

K2 – 1 atom of Potassium has atomic mass of 39.1, but there are 2 atoms here so x2 is 78.2

C  – see on the periodic table, atomic mass 12

O – We already established has atomic mass 16, but there are 3 atoms of it here so x3 is 48

78.2 + 12 + 48 = 138.2

 

Introduction to the mole

The mole is a measure of the amount of a substance.

1 mole of an element is equal to that particular element’s relative formula mass or relative atomic mass. (1mol = RFM and RAM)

Example: 1 mole of oxygen has a molar mass/relative formula mass of 16gmol-1 (the unit gmol-1 literally means grams per mole, so there is 16g per mole of oxygen, 1 mole of oxygen has a mass of 16g)

As we know 16 is oxygen’s relative formula mass, and its relative atomic mass listed on the periodic table.

If we have 2 moles of oxygen, then its ‘molar mass’ (same as relative formula mass) would be 32gmol-1

But what if we are given the weight of an element in grams and told to calculate the amount of moles?

If I have 50g of Oxygen, all I need to do is divide that by Oxygen’s molar mass/relative formula mass (16gmol-1) and I get 3.125 moles. (Test this by multiplying 3.125mol by 16gmol-1).

This is all illustrated in the calculation formula triangle below:

1 mole of a substance contains 6.02×1023 particles/atoms, this number is known as Avogadro’s number and you can’t escape it in chemistry.

A book I highly recommend is Calculations in AS / A Level Chemistry by Jim Clark, goes into detail with moles but keeps it simple and fun, with plenty of practice questions. The book is not just about moles, it’s everything, and has been essential. You can get a second hand version pretty cheap here: http://www.amazon.co.uk/Calculations-AS-A-Level-Chemistry/dp/0582411270/ref=sr_1_cc_1?s=aps&ie=UTF8&qid=1335368574&sr=1-1-catcorr

The bottom row multiplies, the top row divides. So…

moles x molar mass = mass

molar mass x moles = mass (same thing)

mass ÷ moles = molar mass

mass ÷ molar mass = moles

It’s very straightforward, so the above may seem patronising but it’s just in case anyone finds triangles tricky. The other formula triangles are below.





Lesson 2

3 05 2012

Click this link to download: Lesson 2: Atomic Structure, Isotopes and Electronic Configuration

Objective: To understand the basics of atomic structure, Isotopes and electronic configuration.

Atomic Structure

Atoms are made up of sub-atomic particles called protons (+), neutrons (0 hence being neutral) and electrons (-).

Particles Mass Charge
p 1 +1
n 1 0
e 0 -1

Example:

7

   Li

3

In Lithium there are 3 protons, 3 electrons, and 4 neutrons.

The number below the elemental symbol (in this case 3) is the atomic number; this indicates the number of protons in the nucleus of the atom.

The number above (in this case 7) is the atomic mass number; this represents the number of protons + the number of neutrons.

We already know from the atomic number that the number of protons is 3, so if we subtract 3 from 7 we are left with the number of neutrons, 4.

Remember: the number of protons = the number of electrons

(If we assume the element to be electrically neutral, or uncharged. If dealing with questions regarding charge, ions etc, then further considerations would need to be made.)

Test yourself:

39

     K

19

p = ?

n = ?

e = ?

(See answers at the end of this document)

Isotopes

These are atoms of the same element that have the same atomic number, but different mass number.

In other words, Isotopes have the same amount of protons but a differing amount of neutrons.

Example:

35                        37

     Cl                         Cl

17                        17

n = 18                 n = 20

Both Isotopes of Chlorine have 17 protons (and electrons), but as you can see a different amount of neutrons.

Electrons exist around the nucleus of atoms in different energy levels called electron shells, of which there are several.

Sticking with Chlorine, we see that its 17 electrons are spread throughout the layers of electron shells that make up an atom of this element.

The centre is of course the nucleus, where protons and neutrons reside, and the grey dots surrounding it are the electrons.

Those nearest to the nucleus are in the first shell, moving out one is the second shell, and the outer third shell.

In the first shell there are 2 electrons, in the second shell there are 8 electrons and in the third outermost shell there are 7 electrons (hence 2, 8, 7). It is important to know how many electrons each shell can physically hold; as there is a reason every element follows these basic rules, so onto our next heading….

Electronic Configuration

The way in which electrons are arranged within an atom.

Shell

name

Subshell

name

Subshell

max

electrons

Shell

max

electrons

K

1s

2

2

L

2s

2

2+6=8

2p

6

M

3s

2

2+6+10

=18

3p

6

3d

10

N

4s

2

2+6+

+10+14

=32

4p

6

4d

10

4f

14

http://en.wikipedia.org/wiki/Electron_shell

As you can see each shell has a sub-shell, with shell 2 being made up of 2s and 2p, shell 3 being made up of 3s, 3p and 3d. Each sub-shell has a limit Read the rest of this entry »





Lesson 2: Atomic Structure, Isotopes and Electronic Configuration

3 05 2012

Click the link below to download the Lesson, it is a Word Document file.

Lesson 2

If there are any problems downloading, or general questions/thoughts, don’t hesitate to comment.





Chemistry Lesson 1: Periodic Trends

3 05 2012

Click the link to download this lesson: Lesson 1: Periodic Trends

I wrote this to help remember the basics of periodic trends. These are basic but important for science A Level revision.

Click the link at the top and feel free to download the Word Document, leave a comment 🙂

Image courtesy of Wikipedia

Atomic radius

The distance from the atomic nucleus to the outermost stable electron orbital in an atom that is at equilibrium.

↓ Increases down a group (column of the periodic table)

←  Increases across a period (row of the periodic table) from right to left

Therefore atomic radii are largest in the bottom left corner of periodic table.

e.g. Ce is the largest, and Fr has a larger radius than He.

Why does atomic radius decrease as you go across a period? (from left to right) →

The effective nuclear charge increases → therefore attracting the orbiting electrons towards the nucleus and lessening the radius. Less distance between the electrons and the nucleus so the nuclei pull is stronger.

(See bottom of this post for more information on effective nuclear charge.)

Why does atomic radius increase as you go down a group?

The addition of a new energy level (shell)

Ionisation Energy

 This is the energy required to Read the rest of this entry »