Archive for the ‘Berry Reference (Notes)’ Category

STPM Chemistry Form 6 Notes – Terminology and Concepts: Liquid and Solid States (Part 3)

June 23, 2010

This is Part 3 of the series of notes on “Liquid and Solid State” from Berry Berry Easy. In the previous part, the crystal systems were discussed in tabulated form. With that knowledge in hand, it is time to look at the packing of it. It is absolutely important that you draw the following structure at least once and try to have a 3D mental image of it. This action of visualising will help you to understand the information more easily. And when you understand, you don’t have to memorise it anymore, it will come natural to you. (Nonetheless, at least for the beginning, try to memorise them. Subsequent topics will require understanding from this topic, so do not skip this topic)

STPM Form 6 – Terminology and Concepts: Liquid and Solid States (Part 3)

Four types of lattice points:

1. Lattice point at the corner of the unit cell (1/8)
2. Lattice point on the edge of the unit cell (1/4)
3. Lattice point on the face of the unit cell (1/2)
4. Lattice point in the centre of the unit cell (1)

Coordination number – the number of atoms, molecules or ions (called the nearest neighbours) that surrounds a given atom, molecule or ion in a crystal lattice.

A) Simple cubic cell

Example: Caesium chloride & Polonium

• Sphere touches six other spheres.
• Four sphere in its own layer, one sphere above the layer and one sphere below the layer.
• Coordination number = 6
• Unit cell contains in total one atom (8 corners x 1/8 = 1)

B) Body-centre cubic lattice

Example: Sodium, Barium, Potassium, Iron, Manganese, Chromium & Vanadium

• Sphere touches eight other spheres.
• Second layer are placed in the hollows between the spheres in the first layer.
• Each sphere atom is in contact with four atoms in the layer above and four atoms in the layer below.
• Coordination number = 8
• Unit cell contains in total of two lattice points per unit cell (8 corners x 1/8 + 1 = 2)

C) Close-packed structures

Example: Sodium chloride

• Unit cell contains in total of four atoms per unit cell (8 corners x 1/8) + (6 faces x 1/2)

i) Cubic close packing (ABCABCABC) / Face-centered cubic / Simple cubic close packing

Thomas Harriot (1585) first pondered the mathematics of the cannonball arrangement or cannonball stack, which has a face-centered cubic lattice.

• Sphere touches twelve other spheres.
• First layer of spheres is packed as closely and each sphere atom is in contact with six other atoms.
• Second layer of spheres is placed on top of the first layer, so that each sphere in the second layer rests on the hollows between the spheres in the first layer.
• Each sphere atom is in contact with six atoms in its own layer, three spheres (atoms) in the layer above and three spheres (atoms) in the layer below.
• Coordination number = 12

ii) Hexagonal close packing (ABABABABA)

• Sphere touches twelve other spheres.
• First layer and the second layer of spheres are packed in the same way as cubic close packing.
• (Difference = the third layer of spheres is placed on top of the first layer)
• Coordination number = 12

Stay tune to Berry Berry Easy for the next part (Part 4) in this series covering allotropy and enantiotropy.

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SPM Biology Form 4 Notes – Terminology and Concepts: Movement of Substances Across the Plasma Membrane (Part II)

June 20, 2010

Berry Berry Easy presents Part 2 of the SPM Form 4 Biology notes for “Movement of substances across the plasma membrane“. In Part 1, the uniqueness, importance, structure and permeability of plasma membrane were discussed. This part focuses on the gist of the topic, which is on ‘transport’. Both the passive and active transport process is examined in this post. One thing for sure that Berry Readers will remember long after you leave school (yup, 100% sure) is how Osmosis works. It’s a topic which for some unknown reasons will give students headache in the beginning but after you leave school, it all make sense to you. You may ask your elder siblings if they still remember osmosis. Anyway, let the notes begin.

SPM Biology Form 4 – Terminology and Concepts: Movement of Substances Across the Plasma Membrane (Part 2)

Materials must be able to move through the plasma membrane in order for the cell cytoplasma to interact with the external environment. Therefore, the movement of soluble substances can occur in several mechanisms:

• A. Process of Passive Transport
• B. Process of Active Transport

A. Passive Transport

i) Simple Diffusion

• not selective: lipid-soluble molecules, gases and water.
• not control by cell.
• movement of the molecules from a region of higher concentration to a region of lower concentration.
• Factors affecting the rate of diffusion are temperature, size of molecules/ions, diffusion gradient, surface area and diffusion medium.
• example: diffusion of oxygen and carbon dioxide at the alveolus.

ii) Osmosis:

• only water molecules.
• not control by cell.
• movement of water from a region of higher concentration to one of lower concentration and often occurs across a semipermeable membrane.
• strong sucrose solution = less water molecule = low water potential.
• weak sucrose solution = more water molecule = high water potential.
• example: absorption of water by root hairs.

iii) Facilitated Diffusion:

• very specific: glucose, nucleic aicds, amino acids, protein and mineral ions.
• control by cell.
• transport of molecules (only certain molecules) across the outer membrane of living cell by a process of carrier protein (hydrophilic group) / channel protein (Ions: Na⁺, Ca²⁺, K⁺) within the cell membrane.
• normally take place from a region with higher concentration of molecules to a region of lower concentration.
• example: absorption of digested food in the villus.

B. Process of Active Transport

• very specific: minerals ions and amino acids.
• control by cell.
• This process needs carrier proteins and energy (due to against concentration gradient) from a region of lower concentration to a region of higher concentration).
• Cell must expend energy that derived from ATP (adenosine triphosphate)
• example: human nerve cells (sodium ions are constantly transport out of the cell) / ions intake by root hairs of a plant.

Finally, the end of this part. Stay tune for the final part (Part 3) of SPM Biology Form 4 notes on “Movement of substances across the plasma membrane“.

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SPM Biology Form 4 Notes – Terminology and Concepts: Movement of Substances Across the Plasma Membrane (Part I)

June 15, 2010

Berry Berry Easy is back with SPM Form 4 Biology notes for “Movement of substances across the plasma membrane“. The plasma membrane is a membrane of biological nature which forms the barrier between the interior of a cell against external environment. However, do not just think of it as a layer which do not have any function. Typically, movement of substances to and from cells to outer environment is controlled (due to its permeability) by the plasma membrane. The key concept here for this topic is to understand the selective-permeability nature of a plasma membrane. The rest of the topic are rather easy and intuitive. Do note that plasma membrane is also known as cell membrane and plasmalemma.

SPM Biology Form 4 – Terminology and Concepts: Movement of Substances Across the Plasma Membrane (Part 1)

1. Uniqueness of Plasma Membrane (also known as cell membrane):

• it is a semi-permeable cell membrane
• it allows water and certain substances to move in and out of the cell.

2. Importance of Plasma Membrane:

• - cells obtain nutrients and gases
• cells excrete metabolic wastes
• cells can maintain pH for enzyme activity
• cells can maintain ionic concentration of the cells for enzyme activity
• control the types and the amount of substances
• allow useful substance (hormones/enzymes) to secrete from cells
• protect cells
• a boundary between the inside and outside of cell.

3. Structure of the basic unit of plasma membrane

• Phospholipid molecule:
  ‘Head’ – hydrophilic: a polar phosphate molecule (philic~loves water / attracted to water)
  ‘Tail’ – hydrophobic: two non-polar fatty acids (phobic~hates water / repelled to water)
• Formation:
  Hydrophilic heads pointing outwards
  Hydrophobic tails pointing inwards
  (Bilayer phospolipid)

Fluid Mosaic Model (Protein embedded in the bilayer)

Carrier protein

• carrier for some molecules (glucose, amino acids, proteins and nucleic acids)
• controls the movement of ions and particles (Na⁺, Ca²⁺ and K⁺)
• Glycoprotein

Glycolipid

• combination of lipids and polysaccharides

4. Permeability

Permeable (allow to pass through)

• small non-polar molecules (vitamins A, D, E, K, fatty acids, glycerol and steroids)

Impermeable (not allow to pass through but with help of carrier protein and cellular energy, it is allow to pass through)

• large polar molecules (glucose, amino acids, mucleic acids and polysaccharides)
• charged ions (H⁺, Na⁺, K⁺, Cl^- and Ca²⁺)

Substances that are allowed to move in the cell:

• CO₂
• O₂
• excess H₂O
• waste: nitrogenous

Substances that are allowed to move out of the cell:

• CO₂
• O₂
• amino acids
• ionic salts
• glucose

This is the end of Part 1 of this topic. Do check out Part 2 which covers passive and active transport.

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STPM Chemistry Form 6 Notes – Terminology and Concepts: Liquid and Solid States (Part 2)

June 8, 2010

Berry Berry Easy will continue on with the notes about “Liquid and Solid State” with Part 2 in the series. In Part 1, we touched on the changes of state of matter, kinetics theory of liquid and the structure of liquid. This part will be all about crystal lattice. It is another easy portion of the topic as with the previous part. But whenever possible, do try to visualise the crystal lattices and learn to differentiate the characteristics graphically. It is the easiest way to understand this topic. It will also be useful to know more examples of each type of crystal systems than the one being quoted in the notes below. So pick up your pencil and some paper, and let the sketching of crystal lattices begins. (Do refer to books or references online too while you sketch).

STPM Chemistry Form 6 – Terminology and Concepts: Liquid and Solid States (Part 2)

Crystal lattice – regular arrangement of atoms, molecules or ions in a crystalline solid.

Unit cell – a small repeating unit that contains a group of particles (atoms, ions or molecules) in a crystal.

There are 7 crystal systems (primitive unit cells – all the lattice points are placed at the corners of the cell only):

There are 7 multi-primitive lattices (lattice points are located not only at the corners, but also at the faces or centres of the cells).

There are 14 types of crystal lattices.

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STPM Chemistry Form 6 Notes – Terminology and Concepts: Liquid and Solid States (Part 1)

June 3, 2010

Continuing from the previous post about Gas, Berry Berry Easy will be releasing a series of notes about “Liquid and Solid State“. This topic will be longer but should be slightly easier than gas. Most of this topic would seemed as very common and just reinforces what you have already known from daily life. This could be due to the fact that you can ‘feel’ liquid and solid state matter with your fingers and see it more easily than gases. So let us begin the journey to this relatively easy topic. (Note however that, some students still make silly mistakes which meant that they concede easy marks away)

STPM Form 6 – Terminology and Concepts: Liquid and Solid States (Part 1)

Changes in the States of Matter

1. Freezing / Solidification – liquid –> solid
2. Melting – solid –> liquid
3. Evaporation – liquid –> gas / vapour
4. Condensation – gas / vapour –> liquid
5. Sublimation – gas / vapour –> solid
6. Sublimation – solid –> gas / vapour

(Sublimation – iodine, ammonium chloride and solid carbon dioxide)

Kinetics Theory of Liquid

The kinetics energy content of the particles in a liquid is closer to the kinetic energy content of the particles in a solid than to that of a gas.

Important points:

i) Liquid is made up of tiny particles.
ii) Particles in liquid are continually moving in a zigzag.
iii) The motion for particles in liquid are vibration, rotation and translation.
iv) Particles in liquid are not in an orderly arrangement. There are loose clusters of particles which are packed closely.
v) Particles in liquid have strong forces of attraction between the particles.
vi) Particles in liquid have more kinetic energy than the particles in solid but less kinetic energy than particles in gases.

Enthalpy of Fusion – The amount of heat required to change one mole of a pure solid into a liquid.

Enthalpy of Vaporisation – The amount of heat required to change one mole of pure liquid into a gas.

The Structure of a Liquid

i) Melting process:

• Particles move faster when solid is heated.
• The vibrations of the particles increase when temperature of the hot solid increases.
• The particles in the solid acquired sufficient kinetic energy to overcome the attraction forces between particles.
• The particles break away from one another.
• Solid has become liquid.

ii) Freezing process:

• The motion of particles in liquid slows down when liquid is cooled.
• The particles have low kinetic energy.
• The particles in liquid have strong attraction forces between particles to overcome the motion of the particles.
• Particles held in fixed positions in the lattice structure.
• Liquid has become solid.

iii) Vaporisation process (open container that exposed to the atmosphere):

• The particles escape from the surface of the liquid and become gas.
• The rate of vaporisation increases with a rise in temperature, a decrease in external pressure and an increase in the surface area of the liquid.
• A rise in temperature

- room temperature: small percentage of particles have high kinetics and sufficient to overcome the attraction forces between particles and then escape from the surface of the liquid.


A decrease in external pressure (increase in internal pressure)
- particles that have enough kinetics energy to vaporise.
- vapour pressure of liquid increases.
- the particles in liquid collided with one another.
- particles have enough kinetics energy to vaporise.
- a distribution of kinetic energy has formed.


An increase in the surface area of the liquid
- the particles in liquid are collided with one another.
- liquid exposed to the air will evaporate (on top of the liquid).
- particles with higher kinetics energies than the average kinetic energy will escape as gas particles first.

iv) Boiling process:

• Particles move faster when liquid is heated.
• The vibrations of the particles increase when temperature of the hot liquid increases.
• The particles in the solid acquired sufficient kinetic energy to overcome the attraction forces between particles.
• The particles break away from one another.
• Solid has become liquid.

Velocity of the particle ­ increase when

• Temperature ­increase
• Kinetic energy ­increase

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STPM Chemistry Form 6 Notes – Terminology and Concepts: Gas

May 26, 2010

Gas, solid and liquid. (Or some are more familiar with the Malay equivalent of gas, pepejal dan cecair) These are the three classical states of matter. Among the three, liquids and gases can be grouped together to be fluids. So some students get confused when they try to read out of the syllabus. Concentrating back on gases, students typically underestimate the difficulty of “gas” when they study chemistry. Admittedly, it is rather easy to understand gases. One word of caution though from Berry Berry Easy because it is rather hard to master gas, especially gas related calculations. (So pay more attention to it).

At the minimum, students must try to master the concept Ideal Gas Law as soon as possible before they proceed to the calculations. Ideal Gas Law is simple but yet for some unknown reasons, students find it hard to make sense out of it. You will use the Ideal Gas Law in most of the calculations even up to university level for most engineering courses. So for the sake of your STPM and future studies, try and master this simple yet surprisingly difficult (oxymoron) topic of Gas. We’ll kick off with some revision of the states of matter.

STPM Chemistry Form 6 – Terminology and Concepts: Gas

Kinetics Theory of Matter

• describe the behaviour of particles in solids, liquid and gas.

Solid State

• particles are held rigidly in fixed positions by strong attractive forces in an orderly arrangement;
• particles cannot move freely;
• particles can only vibrate or rotate about their mean position;
• particles have less energy (compared to liquids and gases);
• solids cannot be compressed;
• solids have fixed shapes;
• solids have fixed volume

Liquid State

• particles are packed closely together in cluster;
• particles are not in an orderly arrangement;
• particles can vibrate, rotate and move freely;
• particles have more energy (compared to solids) but have less energy (compared to gases);
• liquids are not easily compressed;
• liquids have no fixed shape (take the shape of the container);
• liquids have fixed volume.

Gaseous State

• particles are separated from each other by distance far greater than their own size;
• particles have no forces between the particles.
• particles are not in an orderly arrangement;
• particles can vibrate, rotate and move freely within the container;
• particles have more energy (compared to liquids and solids);
• particles are in constant random motion, moving in straight lines;
• particles collide (elastic) with the walls of the container, they exert a pressure on the container and there is no loss of kinetics energy during the collision;
• gases are easily compressed;
• gases have no fixed shape (take the shape of the container);
• gases have no fixed volume.

Kinetics Theory of Gases

• describe the behaviour of ideal gas.
• the average kinetics energy of gases particles is directly proportional to the absolute temperature of the gas (Kelvin).
• four assumptions associated with this theory:

i.) particles are small compared to the distances between particles that their volumes are negligible.
ii.) particles move in straight lines. The direction of a particle’s motion is changed only by its collision with either another molecule or the walls of the container. All the collisions are to be elastic (no loss of energy).
iii.) particles are in constant random motion. Gas pressure is only caused by collisions of the particles against the walls of the container.
iv.) Gas molecules exhibit no intermolecular forces. The particles neither attract nor repel one another.

Gas Laws

• three common gas laws to know: Avogadro’s Law, Boyle’s Law and Charles’ Law – A, B and C laws of gases.

(If you find yourself about to get confused, here is a simple story about how the scientist, Avogadro might have made his discovery: Avogadro was into counting big numbers, so his law focuses on the number of molecules. Therefore, Avogadro’s law deals with the relationship between moles of gas and volume. Big Boy Boyle sat on his lunch and smashed it (decreased the volume of his sandwich), by increasing the pressure on it. Therefore, Boyle’s law deals with the relationship between pressure and volume. Good ol’ Chuck overheated his popcorn and it scattered all over (increased its volume). Therefore, Charles’s law deals with the relationship between temperature and volume.) taken from General Chemistry Part II Sections VI-X pg 13. (2001) Berkeley Review.

1. Avogadro’s Law

• Amedeo Avogadro (1811)
• equal volumes of all gases at the same temperature and pressure contain equal numbers of molecules.

V / n = k (a constant)
V₁ / n₁ = V₂ / n₂
Where n = number of moles of gas

* Molar volume of a gas (volume occupied by 1 mol of any gas) at standard temperature and pressure (s.t.p.) is 22.4 dm³ (Condition: 0˚C / 273 K and 101.3 kNm^-1 / 1 atm.).

2. Boyle’s Law

• Robert Boyle (1662)
• the volume occupied by fixed mass of gas is inversely proportional to its pressure at constant temperature.
• applies under isothermal conditions in a closed container.

pV = k (a constant)
p₁V₁ = p₂V₂

* Real gases obey Boyle’s law only at low pressures and high temperatures (ideal gas or perfect gas).
* Real gases do not obey Boyle’s law at high pressures and low temperatures (non-ideal behaviour).

3. Charles’ Law

• Jacques Charles (1780)
• the volume occupied by fixed mass of gas is directly proportional to its absolute temperature at constant pressure.

V / T = k (a constant)
V₁ / T₁ = V₂ / T₂

* Temperature is the absolute temperature (-273˚C / 0 K)
* Absolute temperature scale (Kelvin scale) as the temperature -273˚C was adopted as the ‘zero’.

Ideal Gas Equation

Combining Avogrado’s law, Boyle’s law and Charles’ law

• Ideal gas equation:

pV = nRT
where R is a constant and its value of 8.31 J mol^-1 K^-1
pressure: Pa or Nm^-2 (1 atm = 101 kPa)
volume: m³ (1 cm³ = 1 x 10^-6 m³; 1 dm³ = 1 x 10^-3 m³)
temperature: K

n = m / M_r
where m = mass of gas and M_r = relative molecular mass of gas

m / V = ρ
where ρ = density of a gas

4. Dalton’s Law

• the total pressure of a mixture of gases do not react is the sum of the partial pressures of the constituent gases on the mixture.

P_T = P_A + P_B + P_C + …
where  P_T = total pressure of the mixture and

P_A, P_B, P_C = partial pressure of gases A, B and C.
Mole fraction of A (X_A) in a mixture of A and B
= (number of moles of A) / (total number of moles of A + B)
= n_A / (n_A + n_B)

P_A = P_T x X_A
where P_T = total pressure, P_A = partial pressure of gas A, X_A = mole fraction of gas A

5. Deviation from Ideal Behaviour

Factors:

1. pressure
2. temperature
3. molecular size
4. intermolecular forces

Positive deviation (volume of gas molecules):

1. low pressures (molecules are very far apart – volume of the gas molecules by comparison is extremely small and can be ignored)
2. high pressures (molecules are closer together – volume of the gas molecules cannot be ignored)

Negative deviation (intermolecular forces of attraction):

1. low temperature (intermolecular forces of attraction between the molecules will reduce the force exerted by the impact of the molecules collide the wall of container. Pressure exerted by the gas is reduced).
2. high temperature (kinetics energy of the molecules is so high that the intermolecular forces between gas molecules can be ignored).

Negative deviation (polar bonds)

1. Least deviation – hydrogen gas (small molecular size and non-polar. It possesses very weak intermolecular forces of attraction).
2. Marked deviation from ideal behaviour – carbon monoxide gas (polar bonds. It possesses stronger intermolecular forces)

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STPM Chemistry Form 6 Notes – Terminology and Concepts: Matter

May 19, 2010

For some inexplicable reasons, a small percentage of Form 6 students seemed to struggle initially with the concept of matter (The larger percentage tend to understand this intuitively) This is despite being exposed to it in SPM-level physics and chemistry. When the Berry Berry Teacher questioned her previous batches of students why they find it hard to understand the concept of it, they said that since they couldn’t understand it earlier, they have a mental blockage towards it. (Luckily, most students will eventually achieve full understanding on this subject matter.)

Everything you can see in this physical world is consist of substance, generally term as matter. (It’s not an entirely accurate description but technically applicable and correct) Berry Berry Easy advises students to stick to the terminologies used in the syllabus as different fields tend to define matter differently. So to be on the safe side, try and absorb all the terminologies and concepts presented in the notes below.

STPM Form 6 – Terminology and Concepts: Matter

Matter – anything that occupies space and has mass.

Fundamental Particles of Atoms (Historical Point of View)

John Dalton (1808) – atomic theory

1. Atoms – small indivisible particles.
2. Atoms – neither created nor destroyed.
3. Atoms – chemical reactions result from combination / separation of atoms.

J. J. Thomson (1897)

1. Electrons – negatively-charged particles.
2. Atoms – positively-charged sphere.

Ernest Rutherford (1911)

1. Atoms – consists of a positively-charged nucleus with a cloud of electrons surrounding nucleus.
2. Protons – positively-charged particles.

Niels Bohr (1913)

1. Electrons – surrounding the nucleus (orbit).

James Cadwick (1932)

1. Neutrons – electrically neutral subatomic particles.
2. Neutrons – mass almost the same with a proton.
3. Nucleus of an atom – consists of protons and neutrons.

Modern Atomic Model

1. Nucleus of an atom – consists of protons and neutrons.
2. Electrons – moving around the nucleus (orbits / electron shells/ quantum shells)

Atoms

Atom – smallest particle of an element.

Relative atomic mass (A_r) - (an element) average mass of one atom of the element relative to 1/12 times the mass of one atom of carbon-12.
= (average mass of one atom of the element) / (1/12 x mass of one atom of C-12)
Or
= 12 x [(average mass of one atom of the element) / (mass of one atom of C-12)]

Cations – positively-charge ions.
Example: H⁺, K⁺, NH₄⁺ and Mg²⁺

Anions – negatively-charge ions.
Example: Br^-, OH^-, O^2- and S₂O₃^2-

Molecule – a group of two or more atoms.

Relative molecular mass (M_r) – (an element or compound) average mass of one molecule of the substance relative to 1/12 times the mass of one atom of carbon-12.
= (average mass of one molecule of substance) / (1/12 x mass of one atom of C-12)
Or
= 12 x [(average mass of one molecule of substance) / (mass of one atom of C-12)]

Proton number / Atomic number / Number of protons (Z)

• Number of protons in the nucleus of an atom.
• Number of electrons (neutral atom).

Nucleon number / Mass number / Number of nucleon (A)

• total number of protons and neutrons in the nucleus of an atom.
  A = Z + N
  N = number of neutrons

Isotopes (of the same element)

• atoms having the same proton number but different nucleon number.
• same number of protons, number of electrons, electronic configuration and chemical properties.
• different nucleon number, relative mass, density and rate of diffusion.

Relative isotopic mass – the ratio of the mass of one atom of the isotope relative to 1/12 times the mass of one atom of carbon-12 isotope.
= (mass of one atom of the isotope) / (1/12 x mass of one atom of C-12)
Or
= 12 x [(mass of one atom of the isotope) / (mass of one atom of C-12)]

Mass spectrometry

i. Vaporisation chamber – sample is vaporised (produce gaseous atoms or molecules).

ii. Ionisation chamber – vapour is bombarded with a stream of high-energy electrons to form positive ions. X(g) + e –> X⁺(g) + 2e. (produce positive ions)

iii. Acceleration chamber – positive ions are attracted towards the high negative potential plated that accelerates the positive ions to a high and constant velocity. (accelerate the positive ions).

iv. Magnetic Field – accelerated positive ions are deflected into a circular path according to the m/e ratio. (separate positive ions of different m/e ratio)

v. Ion detector – positive ions with different m/e ratios will be deflected to the ion detector that can be recorded on a moving chart. (detect the number and m/e ratio of the positive ions)

vi. Recorder – a flow of current which is amplified and recorded as peaks. (plot the mass spectrum of the sample)

Important note:

• A lighter ion will deflect more than a heavier ion (the same charge)
  Example: ³⁵Cl⁺ will deflect more than ³⁷Cl⁺
• An ion with a higher charge will deflect more than an ion with a lower charge (the same mass)
  Example: ³⁵Cl2⁺ will deflect more than ³⁵Cl⁺

Isotopic abundance = fractional abundance = percentage abundance

One mole – the quantity of a substance that contains the same number of particles (atoms, ions or molecules) as the number of atoms in exactly 12 grams of carbon-12 isotope.

Avogadro constant, L or N_A – number of particles (atoms, ions or molecules) present in a mole of substance (elements or compounds)
= 6.02 x 10²³ (unit is mol^-1)

Number of moles = number of atoms or molecules / Avogadro constant (mol^-1)

Number of particles in a sample = number of moles x Avogadro constant (mol^-1)

Mass (g) = number of moles (n) x M (A_r or M_r)

Number or moles (n) = mass (g) / molar mass (g mol^-1)

Mass (g) = number of moles x molar mass (g mol^-1)

Number of moles = volume of gas (dm³) / 22.4 dm³ at s.t.p. (0˚C and 1 atm or 101 kPa)

Number of moles = volume of gas (dm³) / 24 dm³ at r.t.p. (25˚C and 1 atm or 101 kPa)

Volume of gas (dm³) = number of moles x / 22.4 dm³ at s.t.p.

Volume of gas (dm³) = number of moles x / 24 dm³ at r.t.p.

Number of moles of solute = MV / 1000
(M = concentration in mol dm^-3)
(V = volume in cm³)

Concentration of a solution (g dm^-3) = mass of solute (g) / volume of the solution (dm³)

Concentration of a solution (mol dm^-3) = number or moles of solute (mol) / volume of the solution (dm³)
M_aV_a / M_bV_b = a/b
M₁V₁ = M₂V₂

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SPM Chemistry Form 5 Notes – Terminology and Concepts: Carbon Compounds (Part 9 – Final)

May 14, 2010

Finally, Berry Berry Easy will be ending the long running Carbon Compound series for SPM Form 5 Chemistry with Part 9. This final part won’t be a recapitulation of what have been learnt but rather focuses on polymers. One of the most important natural polymers is natural rubber. And since rubber is historically an important export for Malaysia, it would be beneficial to know about rubber in general. So we present you with the final part of the long series.

SPM Form 5 – Terminology and Concepts: Carbon Compounds (Part 9 – Final)

Polymers

1. Polymer – many small units (monomers) joining together to formed large molecule.

2. Polymer can be classified into two groups:

• synthetic polymers / man-made polymers (polythene; PVC – polyvinyl chloride; artificial silk; and polypropene)
• natural polymers (natural rubber; starch; cellulose; and proteins)

3. Natural polymer: Carbohydrates (polysaccharides) (starch, glycogen and cellulose)

• General formula: C_x(H₂O)_y with the ratio of H:O = 2:1
• Carbohydrates have cyclic structure.
• Monomer: glucose (C₆H₁₂H₆)
• Reaction to form polymer: condensation reaction (- H₂O)

4. Natural polymer: Protein (polypeptide)

• Protein consists of carbon, hydrogen, oxygen and nitrogen (some have sulphur, phosphorus and other elements)
• Monomer: amino acids
• Amino acids have two functional group which are carboxyl group (-COOH) and amino group (-NH₂)
• Reaction to form polymer: condensation reaction (- H₂O)

5. Natural polymer: Natural rubber

• Extracted from the latex of rubber tree (Hevea brasiliensis) which the tree originates from Brazil.
• A molecule of rubber contains 5000 isoprene units.
• Monomer: isoprene, C₃H₈ or 2-methylbuta-1,3-diene.
• Reaction to form polymer: additional polymerisation (one of the double bond in isoprene becomes single bond)

6. Structure of rubber molecule

• Latex is colloid (35% rubber particles and 65% water).
• Rubber particle contains rubber molecules which are wrapped by a layer of negatively-charged protein membrane. Same charge of rubber molecules repels each other. This prevent rubber from coagulate.

7. Coagulation process of latex

The process for the coagulation of latex is summarised as:

1. Acid (H⁺) can neutralise the negatively-charged protein membrane. Example of acid: formic acid, methanoic acid, suphuric acid and hydrochloric acid.
2. The rubber molecules will collide after the protein membrane is broken.
3. Rubber molecules (polymers) are set free
4. Rubber molecules combine with one another (coagulation).

8. Natural coagulation process of latex

For the natural coagulation of latex:

1. Latex is exposed to air without adding acid (duration – overnight).
2. Coagulation process occurs in slower pace due to the bacteria (microorganism) action which produce acid)

9. Prevent coagulation process of latex

The following are latex coagulation prevention method:

1. Alkaline / Basic solution is added to the latex. Example: ammonia (NH₃).
2. Positively-charged hydrogen ion / H⁺ produced by bacteria can be neutralised by negatively-charged hydroxide ion / OH^- from ammonia solution.

10. Properties of natural rubber

• elastic
• cannot withstand heat (become sticky and soft – above 50°C; decompose – above 200°C; hard and brittle – cooled)
• easily oxidised (present of C=C)
• insoluble in water (due to the long hydrocarbon chains)
• soluble in organic solvent (propanone, benzene, petrol etc.)

11. Vulcanisation of rubber

Vulcanisation – process of hardening rubber and increases rubber elasticity by heating it with sulphur or sulphur compounds.

Methods:

• heating natural rubber with sulphur at 140°C using zinc oxide as catalyst or
• dipping natural rubber in a solution of disulphur dichloride (S₂Cl₂) in methylbenzene.

12. Properties of vulcanisation of rubber

• The sulphur atoms are added to double bonds in the natural rubber molecules to form disulphide linkages (-C-S-S-C-) / sulphur cross-links between the long polymer chains. Therefore, vulcanised rubber is more elastics and stronger.
• This increases the molecular size and the intermolecular forces of attraction between rubber molecules. Therefore, vulcanised rubber is more resistant to heat (does not become soft and sticky when hot).
• This also reduces the number of carbon-carbon double bonds in rubber molecules. Therefore, vulcanised rubber is more resistant to oxygen, ozone, sunlight and other chemicals.

13. Comparison between the properties of vulcanised rubber and unvulcanised rubber

14. R & D of rubber

• RRIM – Rubber Research Institute of Malaysia
• MRB – Malaysian Rubber Board
• Rubber Technology Centre
• Various local higher institutions of learning

So finally you are done with the chapter of Carbon Compound if you have mastered all nine parts which was published on Berry Berry Easy.

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