CHEMISTRY F4 F5

Chemistry Form 4 02 Structure of Atom I | Answer 02 Structure of Atom II | Answer 02 Structure of Atom III 02 Structure of Atom IV 05 Chemical Bond I 05 Chemical Bond II 05 Chemical Bond III 05 Chemical Bond IV 06 Electrolyte (Card) 07 Colour of Ion - Copper (Card) 07 Colour of Ion - Iron (Card)	(iPaper) (iPaper) (iPaper) (iPaper)
Chemistry Form 5 Rate of Reaction I Rate of Reaction II Carbon Compound I Carbon Compound II
Complete Formulae List (Highly recomended) Chemical Reaction List - Form 4 Chemical Reaction List - Form 5

 

 

Kimia Tingkatan 4 dan 5

• Latihan Jadual Berkala 2
• Latihan Jadual Berkala 1
• Chemical Equation

Form 4 Chapter 4 - Periodic Table of Elements

Same disclaimer as usual, refer to previous posts. You should read up on the experiments in this chapter since I'll not cover them here. Maybe next time.

Development of the Periodical Table (Paper 1):
Antoine Lavoiser
- classify elements into four groups
- inaccurate - light and heat included

Johann W. Dobereiner
- groups of three called triads
- did not apply to most of other elements

John Newlands
- increasing nucleon number - law of octaves
- only accurate for first 17 elements
- no positions allocated for elements yet to be discovered

Lothar Meyer
- similar chemical properties - volume of atoms of elements
- properties of elements recur periodically

Dmitri Mendeleev
- increasing atomic mass
- gaps for undiscovered elements

Henry Moseley
- increasing proton number/atomic number
- predicted the existence of undiscovered elements

Modern Periodic Table:
(Note: This part comes out a lot in my exams, maybe it'll come out in SPM?)
(Paper 1, Paper 2)
Group 1 elements - Alkali metals
Group 2 elements - Alkaline earth metals
Group 17 elements - Halogens
Group 18 elements - Inert gases

Group 18 elements:
Name: rare / inert / noble gases
Properties: mono-atoms
Down the group:
- Atomic radius increases
- Melting and boiling point increases
- Density increases
Additional info: No reaction, stable element.

Group 1 elements:
Name: alkali metals
Properties: soft solid
Down the group:
- Atomic radius increases
- Melting point decreases
- Density increases
- Electropositivity (Tendency to lose electrons) increases
- Reactivity increases

Group 17 elements:
Name: halogens
Properties: diatomic molecules
Down the group:
- Atomic radius increases
- Melting and boiling point increases
- Density increases
- Electronegativity decreases
- Reactivity decreases

Period 3 elements:
Across the group:
- Atomic radius decreases
- Melting point and boiling point increases and the decreases
- Density decreases
- Electronegativity increases

Explanations:
(Note that covalent compounds experience Van der Waals' forces of attraction (VW). )
Group 18
1. Melting point & boiling point (m.p & b.p)
* From He to Ra, m.p & b.p increases.
- proton number increases
- distance from nucleus to electrovalence increases
- radius increases
- size increases
- VW increases
- more energy needed to weaken VW
- m.p & b.p increases

2. Density
* From He to Ra, density increases.
- similar as m.p & b.p (follow the colour)

Group 1
1. Atomic size
* Increases from Li to Fr
- proton number increases
- orbital number increases
- distance from nucleus to outer shell increases
- size increases

2. m.p & b.p
* Decreases from Li to Fr
- similar
- forces of attraction in metals are getting weaker
- less energy is used to weaken the attraction
- m.p & b.p decreases

3. Reactivity
* Increases from Li to Fr
- similar
- attraction decreases
- Group 1 has 1 electron valence
- easier to let go of electron in outer shell
- reactivity increases

Form 4 Chapter 4 - Periodic Table of Elements 2

Group 17
1. Atomic Size
* Increases from F to At
- proton number increases
- distance from nucleus to electrovalence increases
- radius increases
- size increases

2. Density
* Increases from F to At
- same as above

3. Melting point and boiling point (m.p + b.p)
* Increases from F to At
- same as above
- Van der Waals' forces of attraction (VW) increases
- more energy is needed to weaken VW
- m.p + b.p increases

4. Reactivity
* Decreases from F to At
- proton number increases
- distance from nucleus to outer shell is further
- forces of attraction is weaker
- reactivity decreases

5. Electronegativity
* Decreases from F to At
- the most reactive is the most electronegative (Fl)

Period 3 Elements
1. Atomic Size
* Decreases from left to right
The atomic size does not increase as we move from the left to the right of the periodical table. (This is a common error made by students.)
- proton number increases
- forces of attraction increases
- this pulls the external shell of electrons closer to the nucleus, making the size of the atom smaller.

2. m.p + b.p
* Metal > Gas
- Metals have a stronger bond between atoms - more heat needed to break the bonds
- Gases have a weaker bond between atoms - less heat needed compared to metals

3. Density
* Decreases from left to right
- Metals - atoms are arranged closely to each other (atomic radius increases) making it more dense
- Gases - atoms move around freely due to a weak VW, making it less dense compared to metals

Form 4 Chapter 4 - Periodic Table of Elements 3

Form 4 Chapter 4 - Periodic Table of Elements 3

Final part for this chapter:

More about Periods in the Periodic Table:
Reactions:
1) metal oxide + water --> basic solution
2) non-metal oxide + water --> acidic solution
3) oxide + acidic/basic solution will produce different results, as seen below.
- metal oxide + acid --> salt + water
- metal oxide + base --> no reaction
- non-metal oxide + acid --> no reaction
- non-metal oxide + base --> salt + water
- amphoteric + acid --> salt + water
- amphoteric + basic --> salt + water
(An example of an amphoteric substance is aluminium oxide, Al(OH)3.)

Transition metals

Characteristics:
- high b.p & m.p
- high density
- conductors
- shiny
- ductile
- malleable
- can form a coloured solution in a compound
Examples:
Fe 2+ ions - Green
Fe 3+ ions - Brown
Cu 2+ (Copper) ions - Blue
Co 2+ (Cobalt) ions - Pink
CrO4 2+ (Chromate) ions - Yellow

- catalyst in a reaction
Examples:
Note that <--X--> indicates a reversible reaction with transition metal X.
a) N2 + 3H2 <--Fe--> 2NH3 (Found in the Haber Process)
b) Zn + H2SO4 --CuSO4--> ZnSO4 + H2
c) Contact Process
- S + O2 --> SO2
- 2SO2 + O2 <--V2O5--> 2SO3 (Vanadium Pentoxide / Vanadium (V) oxide)
- SO3 + H2SO4 --> H2S2O7
- H2S2O7 + H2O --> 2H2SO4
Note that the Contact Process involves the production of sulphuric acid. (Chapter 9)

- forms complex ions/compounds (KMnO4 - potassium permanganate, K2Cr2O7 - potassium dichromate, K4Fe(CN)6 - potassium hexacyanoferrate(II).)
- compounds of transition metals contains more than one oxidation number
Example: Iron(II) ions, Iron(III) ions, Copper(I) ions, Copper(II) ions.

Form 4 Chapter 5 - Chemical Bonds

Form 4 Chapter 5 - Chemical Bonds

A relatively short chapter, usually asked in Paper 2 Section B/C.

Definitions:
Ionic compound - Chemical compound formed by the process of releasing and gaining electrons between metals and non-metals.
Covalent compound - Chemical compound formed by the sharing of electrons by a non-metal with a non-metal.
Ionic bond - A strong bond formed when positive ions and negative ions are attracted by the electrostatic forces.
Covalent bond - A weak bond formed when electrons are attracted by the Van der Waals' forces.

How to answer questions about...
Formation of ionic compounds:
- list the electron configuration
- electron valence
- needs to gain/lose electron to form an octet structure
- atom will donate/take electrons to form an octet structure
- valence electron of atom is X, so atom needs to gain/lose electrons to form an octet structure
- when Y atom loses/gains X electrons, it will form an ion in the charge of +/- X (Repeat for other element)
- the + ion and the - ion charge are attracted to each other and form an electrostatic forces
- this will form an ionic bond between ions and form Z (name the compound)
- draw the diagram

Formation of covalent compounds:
- list the electron configuration
- each X atom needs Y electrons to form an octet structure
- non-metal - each X atom will share Y pair of electrons with another X atom
- this will form a Van der Waals' force and later a single/double covalent bond
- draw the diagram

Properties of ionic compounds:
- dissolves in water (not all)
- does not dissolve in organic solvent
- conducts electricity in molten or aqueous solution
--> in ionic compounds, when in molten or aqueous solution, all ions are moving freely, hence it conducts electricity.
- all ionic compounds have a very high m.p & b.p
--> when +ve and -ve charge attracts each other, this forms a strong bond (electrostatic forces)
--> in an ionic compound there is a strong electrostatic force between the +ve and -ve ions
--> a lot of heat and energy is needed to weaken the forces

Properties of covalent compounds:
- dissolves in organic solvent
- does not dissolve in water
- all covalent compounds have a very low m.p & b.p
--> in covalent compounds, there is a very weak VdW force between +ve and -ve ions
--> little heat and energy is needed to weaken the forces
- all covalent compounds do not conduct electricity in any state
--> in covalent compounds no ions exist because all covalent compounds are molecules

That ends the short but important chapter.

Chemical Bonding for SPM Form 4 Chemistry

Covalent bond is the second type of chemical bonds that . Covalent bonds occurs when electrons are shared between atoms, rather than a complete transfer of electrons in ionic bonding. Typically, covalent bonds occur for non-metals when they bind together due to similar tendency for electrons (usually to gain electrons in the syllabus). When non-metals gain electrons, they will share electrons in order to fill up their valence shell, with one of the simplest example being the abundant hydrogen gas. Now that you know that covalent bonding is usually with non-metals and it is different from ionic bonding.
This Part 3 focuses on covalent bonds (definition), non-metals (and how they can form covalent bonds), covalent compound formula prediction, structure of covalent compounds and some tips on covalent bonding. So take note of the notes from Part 2 (Ionic bonding) and try to find parallels from it. This way you’ll better learn the two tips of bonds.
[Tips: In event that you forget how covalent bond works, think of the humble and simplest hydrogen gas. Why is it a good example, because it is easy to remember. As hydrogen atom (H) each has one valence electron in their first electron shell, they will 'prefer' to have a second electron to fill up the first electron shell. If you remember, the first electron shell has a capacity of 2 electrons. When the first H atom wants a second electron, another H atom also wants the same thing. As such, both hydrogen atoms will come together (or in chemistry, we call it 'react') to form H₂, a gas compound. Thus, both atoms now 'enjoy' the stability afforded by full valence shell. So take out your pencil and draw the concept of covalent bonding a few times with the hydrogen atoms.]

Molecules

Covalent Bonds

• It is a chemical bond formed from the sharing of valence electrons between non-metal atoms to achieve the stable duplet of octet electron arrangement.
• Each shared pair of electrons is as one covalent bond.
• It produces molecules.
• Usually the covalent bonds form between non-metal atoms from Group 15, 16 and 17 and sometimes can be formed from Group 14 (carbon and silicon) and hydrogen.
• Covalent bond can be formed from atoms of the same element and atoms of different elements.

Example:

Non-metal	+	Non-metal	–>	Covalent compound
Bromine	+	bromine	–>	Bromine (Br2)
Nitrogen	+	nitrogen	–>	Nitrogen (N2)
Carbon	+	chlorine	–>	Tetrachloromethane (CCl4)
Hydrogen	+	oxygen	–>	Water (H2O)
Hydrogen	+	nitrogen	–>	Ammonia (NH3)

Types of covalent bond formed:

• Single bond = one pair of electrons shared between two atoms.
• Double bond = two pair of electrons shared between two atoms.
• Triple bond = three pair of electrons shared between two atoms.

Non-metal
Group 15

• A nitrogen atom with an electron arrangement of 2.5 needs three more electrons to achieve stable octet electron arrangement after it contribute (through sharing) three valence electrons to another atom (can be from Group 14, 15, 16, 17).
• A phosphorus atom with an electron arrangement of 2.8.5 need three more electrons to achieve stable octet electron arrangement after it contribute (through sharing) three valence electrons to another atom (can be from Group 14, 15, 16, 17).

Group 16

• An oxygen atom with an electron arrangement of 2.6 needs two more electrons to achieve stable octet electron arrangement after it contribute (through sharing) two valence electrons to another atom (can be from Group 14, 15, 16, 17).
• A sulphur atom with an electron arrangement of 2.8.6 need two more electrons to achieve stable octet electron arrangement after it contribute (through sharing) two valence electrons to another atom (can be from Group 14, 15, 16, 17).

Group 17

• A fluorine atom with an electron arrangement of 2.7 needs one more electron to achieve stable octet electron arrangement after it contribute (through sharing) one valence electron to another atom (can be from Group 14, 15, 16, 17).
• A chlorine atom with an electron arrangement of 2.8.7 need one more electron to achieve stable octet electron arrangement after it contribute (through sharing) one valence electron to another atom (can be from Group 14, 15, 16, 17).

Predict the Formula of a Covalent Compound

• Non-metal X atom (valence electron is a)
• Combine with another non-metal Y atom (valence electron is b)
• b = simplest ratio (n) and a = simplest ratio (m)
• Formula of a covalent compound formed, X_nY_m

Example:
The electron arrangement of atom X is 2.8.6 and atom Y has four valence electrons. Which of the following is the formula of the compound formed between X and Y?
(A) Y₄X
(B) Y₂X
(C) YX
(D) YX₂
Solution:

• X has 6 valence electrons, it needs to share 2 electrons to achieve the stable octet electron arrangement.
• Y has 4 valence electrons, it needs to share 4 electrons to achieve the stable octet electron arrangement.
• Therefore, the formula of the covalent compound is X₄Y₂ = Y₂X₄ = simplest ratio YX₂.

Answer: D
Some common covalent compound

• Hydrogen molecule, H₂ (single bond)
• Chlorine molecule, Cl₂ (single bond)
• Bromine molecule, Br₂ (single bond)
• Fluorine molecule, F₂ (single bond)
• Water molecule, H₂O (single bond)
• Nitrogen trifluoride molecule, NF₃ (single bond)
• Tetrachoromethane / carbon tetrachloride, CCl₄ (single bond)
• Ammonia molecule, NH₃ (single bond)
• Oxygen molecule, O₂ (double bond)
• Carbon dioxide molecule, CO₂ (double bond)
• Nitrogen molecule, N₂ (triple bond)
• Ethyne molecule, C₂H₂ (triple bond)

Structure of covalent compounds

• Can be simple molecular structure or giant molecular structure.
• The atoms in the molecule are joined together by strong covalent bond but intermolecular forces are weak by weak van der Waals’ forces.

Berry Important Notes:
In the diagram of ionic compound, always shows

• The outermost shells of all atoms must achieve a stable duplet or octet electron arrangement through sharing.
• The outermost shells of each atom must overlap.
• Label all atoms clearly.

Chemical Bonding for SPM Form 4 Chemistry will be on the differences between ionic compound and covalent compound (in terms of particle, electrons, forces, state, melting point, volatility, solubility in water, solubility in organic solvent, electricity conductor) and also the uses of covalent compound as solvent. So stay tuned.

CHEMISTRY SPM-Nota Tambahan

CHEMISTRY FORM 4

CHAPTER 3: CHEMICAL FORMULAE AND EQUATIONS
3.1 Understanding of mole concept
3.2 Mass and moles of an element
3.3 Empirical formula of magnesium oxide
3.4 Empirical formula of copper (II) oxide

CHAPTER 4: PERIODIC TABLE OF ELEMENTS
4.1 Introduction of periodic table
4.2 Physical properties of Group 1
4.3 Reactivity of Group 1
4.4 Physical properties of Group 17
4.5 Reactivity of Group 17
4.6 Characteristics of transition metal

CHAPTER 5: CHEMICAL BONDS
5.1 Ionic bond
5.2 Covalent bond
5.3 Lewis dot structure of covalent bond
5.4 Electrical conductivity of ionic compound

CHAPTER 6: ELECTROCHEMISTRY
6.1 Understanding of electrolyte
6.2 Understanding of electrolysis
6.3 Electrolysis of concentrated sodium chloride
6.4 Types of electrodes affect product of electrolysis
6.5 Electrolysis of copper sulphate using carbon and copper
6.6 Electroplating of metal
6.7 Purification of metal
6.8 Extraction of metal
6.9 Daniell cell
6.10 Displacement of metal from salt solution
6.11 Constructing electrochemical series

CHAPTER 7: ACIDS AND BASES
7.1 Understanding of acid and base
7.2 Strength of acid
7.3 Concentration of solution - molarity
7.4 Titration method

CHAPTER 8: SALTS
8.1 Preparation of insoluble salt
8.2 Application of precipitation of salt

CHAPTER 9: MANUFACTURED SUBSTANCES IN INDUSTRY
9.1 Contact process


CHEMISTRY FORM 5

CHAPTER 1: RATE OF REACTION
1.1 Understanding of rate of reaction
1.2 Collision theory
1.3 Catalyst affects the rate of reaction

CHAPTER 2: CARBON COMPOUNDS
2.1 Differences between alkanes and alkenes
2.2 Manufacture of ethanol
2.3 Carboxylic acid

CHAPTER 3: OXIDATION AND REDUCTION
3.1 Redox reaction
3.2 Rules of oxidation number
3.3 Redox reaction in displacement of metal
3.4 Rusting of iron
3.5 Redox reaction in electrochemistry reaction