o level chemistry notes

Key things to study for each topic in O levels Disclaimer: This is meant as a summary so not every thing that could possibly come out in your exams may be included. This is also a work in progress so there may be errors or items that are left out by accident. Atomic structure Structure of atom Subatomic particle Symbol Relative charge Relative mass Proton p +1 1 Neutron n 0 1 electron e –1 Isotopes Definition of isotopes: Isotopes are atoms of the s ame element with the same number of protons, but different number of neutrons Calculation of relative atomic mass Q: If chlorine has 2 isotopes with relative masses of 35 and 37 and relative abundance of 75% and 25% respectively, calculate the relative atomic mass of chlorine. A: Relative atomic mass = 37 x 0.25 + 35 x 0.75 = 35.5 Calculation of relative abundance Q: If boron has 2 isotopes with relative masses of 10 and 11, given that boron has a relative atomic mass of 10.8, calculate the relative abundance of boron-10 and boron-11. A: Let the relative abundance of boron-10 be x 10x + 11(1 – x) = 10.8 x = 0.2 Therefore relative abundance of boron-10 is 20% and relative abundance of boron-11 is 80%.

Chemical bonding Template for answering common chemical bonding questions General template 1. State the structure 2. Describe the structure 3. Relate the relevant part of the structure to the physical property specified in the question Structure Simple covalent Giant covalent Giant Ionic lattice Giant Metallic Melting and boiling points • (substance) has a simple covalent structure. • It consists of discrete molecules held together by weak Van der Waals’ forces of attraction. • Very little energy is required to overcome these forces and (substance) has low melting and boiling point. Diamond • Diamond has a giant covalent structure • Each carbon atom is covalently bonded to 4 other carbon atoms in a giant covalent network. • These strong covalent bonds require large amount of energy to overcome. Graphite • Graphite has a giant covalent structure. • Each carbon atom is covalently bonded to 3 other carbon atoms in a giant covalent network. • These strong covalent bonds require large amount of energy to overcome. • (substance) has a giant ionic structure. • It is held together by strong electrostatic forces of attraction between oppositely charged ions in a giant ionic lattice. • A lot of energy is required to overcome these forces and (substance) has high melting and boiling point. • (substance) has a giant metallic structure. • It is held together by strong electrostatic forces of attraction between the positively charged cations and the ‘sea of delocalized’ electrons • A lot of energy to overcome those forces. Hence it has a high melting and boiling point.

Electrical conductivity • (substance) has a simple covalent structure. • It consists of molecules that are uncharged/ electrically neutral. • Since it does not contain any mobile ions or electrons/ or mobile charged particles, it cannot conduct electricity. Diamond • Diamond has a giant covalent structure. • Each carbon atom uses up all four of its valence electrons to form covalent bonds with four other carbon atoms to form a giant covalent network. • It does not have any mobile electrons to conduct electricity. Graphite • Graphite has a giant covalent structure. • Each carbon atom uses 3 out of 4 of its valence electrons to form covalent bonds with 3 other carbon atoms in a giant covalent network. • Hence each carbon atom has one mobile valence electron that is delocalized between the layers to conduct electricity. • (substance) has a giant ionic structure. • In solid state , the oppositely charged ions can only vibrate about their fixed positions due to the strong electrostatic forces of attraction. • Therefore the ions are not mobile to conduct electricity. • (substance) has a giant ionic structure. • In the aqueous/molten state , the electrostatic forces between oppositely charged ions are weakened. • Hence, the ions become mobile and are able to conduct electricity. • (substance) has a giant metallic structure. • It is held together by electrostatic forces of attraction between the positively charged cations and the ‘sea of delocalized’ electrons. • Hence (substance) has mobile electrons that allow it to conduct electricity.

Hardness NA Diamond (same explanation as mp and bp) • Diamond has a giant covalent structure. • Each carbon atom is covalently bonded to 4 other carbon atoms in a giant covalent network. • These covalent bonds require large amount of force to overcome. Graphite • Graphite has a giant covalent structure. • Each carbon atom is covalently bonded to 3 other carbon atoms in a giant covalent network. • Graphite is soft and slippery because the layers of carbon atoms are held by weak Van der Waals’ forces of attraction and can slide over each other. (same explanation as for mp and bp) • (substance) has a giant ionic structure. • It is held together by strong electrostatic forces of attraction between oppositely charged ions in a giant ionic lattice, • which require a lot of force is to overcome. • Pure metals are malleable/ ductile because the orderly layers of same-sized atoms can slide easily over each other without disrupting the structure. • Alloys are harder and stronger than pure metals because alloys have atoms of different sizes. This disrupts the orderly arrangement of atoms and prevent the layers of atoms from sliding over each other easily when a force is applied. Uses of diamond and graphite Substance Uses Related property Diamond Used as drill bit or cutting tool Hard Graphite Used as lubricant and in pencil lead Used as electrodes in electrolysis Soft Conducts electricity

Strength of ionic bonds • The higher the charge of the ions, the greater the strength of ionic bonds. • The smaller the size of the ions, the greater the strength of ionic bonds. Example: Why is the melting point of magnesium fluoride hig her than that of sodium fluoride? • Mg 2+ has a higher charge than Na +. • Mg 2+ is smaller than Na+. • Therefore Mg 2+ has a higher charge density than Na +. • More energy is needs to overcome stronger electros tatic forces of attraction between oppositely charged ions. Strength of metallic bonds • The higher the number of delocalised electrons, the greater the strength of metallic bonds. • The higher the charge of metallic cation, the greater the strength of metallic bonds. • The smaller the size of metallic cation, the greater the strength of metallic bonds. Example: Why is the melting point of magnesium higher than that of sodim? • Mg has more valence electrons and therefore more d elocalised electrons than sodium. • Mg 2+ has a higher charge than Na +. • Mg 2+ is smaller than Na+. • Therefore Mg 2+ has a higher charge density than Na +. • More energy is needs to overcome stronger electros tatic forces of attraction between metallic cations and sea of delocalised electrons. Hydrogen bonding Criteria for hydrogen bonding 1) Hydrogen atom covalently bonded to O, N or F 2) An O, N or F atom with available lone pair(s) Effects of hydrogen bonding 1) Higher boiling point compared to most other simple molecules of similar molecular mass. 2) Density of ice is less than density of water. Example: Why is the boiling point of water higher than methane, CH 4? • Water contains hydrogen bonding between water mole cules. • Methane has Van der Waals’ forces between methane molecules. • Hydrogen bonding is stronger than Van der Waals’ f orces, so more energy is required to overcome the hydrogen bonding in water.

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