Chemistry Guru Podcast

Singapore's top JC Chemistry Tutor Maverick Puah shares his knowledge via his video lessons which follow the new H2 Chemistry syllabus (Subject Code 9729) closely. Students taking A Level Chemistry be sure to subscribe!

How to deduce Nature of Salt in Salt Hydrolysis 11/18/2018

How to deduce Nature of Salt in Salt Hydrolysis A salt can be acidic, neutral or alkaline. We can deduce the nature of the salt from its constituent ions, namely: 1. ion formed from a weak acid will be a conjugate base, eg CH3COO- (from weak acid CH3COOH) 2. ion formed from a strong acid will be neutral, eg Cl- (from strong acid HCl) 3. ion formed from a weak base will be a conjugate acid, eg NH4+ (from weak base NH3) 4. ion formed from a strong base will be neutral, eg Na+ (from strong base NaOH) This can be summarised as shown: salt hydrolysis nature of ions summary Next we can move on to deduce the nature of a salt. Let's have some examples. 1. Sodium chloride NaCl is neutral What we do is simply just deduce the nature of each constituent ion based on the acid or base that forms that ion. salt hydrolysis sodium chloride neutral Na+ is neutral since it is formed from strong base NaOH Cl- is neutral since it is formed from strong acid HCl Since both cation and anion are neutral, NaCl must be neutral. 2. Sodium ethanoate CH3COONa is alkaline salt hydrolysis sodium ethanoate is alkaline Na+ is neutral since it is formed from strong base NaOH. CH3COO- is alkaline since it is the conjugate base of weak acid CH3COOH. Since we have a conjugate base, CH3COO- will dissociate in water to give OH-, hence the salt is alkaline. 3. Ammonium chloride NH4Cl is acidic salt hydrolysis ammonium chloride is acidic Cl- is neutral since it is formed from strong acid HCl. NH4+ is acidic since it is the conjugate acid of weak base NH3. Since we have a conjugate acid, NH4+ will dissociate in water to give H+, hence the salt is acidic. Watch this video to learn an easy way to understand salt hydrolysis, a concept that students often find confusing in Ionic Equilibria! Topic: Ionic Equilibria, Physical Chemistry, A Level Chemistry, Singapore You can also view this with screenshots and detailed explanation Do check out the following for more video lessons: If you are looking for H2 Chemistry Tuition, do consider taking up my . More info at You can also find out more about my /episode/index/show/alevelchemistrylesson/id/7600028

Bronsted Acids and Bases: Types and Strength 11/11/2018

Bronsted Acids and Bases: Types and Strength According to Bronsted-Lowry theory of acids and bases, - an acid is a proton, H3O+ or H+ donor - a base is a proton, H3O+ or H+ acceptor The strength of an acid or base is related to the extent of dissociation in solution: - a strong acid or base is fully dissociated - a weak acid or base is partially dissociated So therefore we will have 4 possible types of acids and bases: 1. Weak Acid 2. Strong Acid 3. Weak Base 4. Strong Base Also, each of these acids and bases will dissociate in solution to give different species. 1. Weak Acid Dissociation eg CH3COOH bronsted theory weak acid CH3COOH dissociation CH3COOH is a weak acid and dissociates partially in solution (as indicated with reversible arrow) to form H+ and CH3COO- ions. Since this is a reversible process, CH3COO- can accept H+ to form back CH3COOH. Therefore the nature of CH3COO- is basic and we call CH3COO- the conjugate base of CH3COOH. 2. Strong Acid Dissociation eg HCl bronsted theory strong acid HCl dissociation HCl is a strong acid and will dissociate fully (as indicated with full arrow) in solution to form H+ and Cl- ions. Since this is an irreversible process, Cl- has no tendency to accept H+ and form back HCl. Therefore the nature of Cl- is neutral. 3. Weak base dissociation eg NH3 bronsted theory weak base NH3 dissociation NH3 is a weak base and dissociates partially in solution (as indicated with reversible arrow) to form OH- and NH4+ ions. Since this is a reversible process, NH4+ can donate H+ to form back NH3. Therefore the nature of NH4+ is acidic and we call NH4+ the conjugate acid of NH3. 4. Strong base dissociation eg NaOH bronsted theory strong base NaOH dissociation NaOH is a strong base and will dissociate fully (as indicated with full arrow) in solution to form OH- and Na+ ions. Since this is an irreversible process, Na+ has no tendency to donate H+ and form back NaOH. Therefore the nature of Na+ is neutral. So in summary: - weak acid will dissociate to give conjugate base - weak base will dissociate to give conjugate acid - strong acid will dissociate to give a neutral ion - strong base will dissociate to give a neutral ion bronsted theory dissociation summary These concepts are fundamental and very important in Ionic Equilibria. Check out this video lesson to learn more about acids and bases and their dissociation in solution! Topic: Ionic Equilibria, Physical Chemistry, JC, H2, A Level Chemistry, Singapore You can also view this with screenshots and detailed explanation. Do check out the following for more video lessons: If you are looking for H2 Chemistry Tuition, do consider taking up my . /episode/index/show/alevelchemistrylesson/id/7515992

Balance Redox Reaction via Half Eqn Method 11/11/2018

Balance Redox Reaction via Half Eqn Method Balancing Redox Equations via the Half-Equation Method can be done via the following systematic steps. 1. Balance Redox Equation in Acidic Medium Let's have this reaction as an example: balance redox reaction using half equation method example We can determine which species is oxidised and reduced by comparing the oxidation number: - Oxidation state of manganese decreases from +7 to +4 hence MnO4- is reduced - Oxidation state of iodine increases from -1 to 0 hence I- is oxidised We can then write down the half equations for oxidation and reduction: balance redox reaction half equations Next we need to balance each half equation. For half equation in acidic medium, the steps are: 1. Balance elements oxidised or reduced 2. Balance oxygen using water 3. Balance hydrogen using H+ 4. Balance charge using electron So after applying these 4 steps for each half equation, we'll end up with these two balanced half equations: balance redox reaction half equation balanced in acidic medium Next we need to modify each half equation so that the number of electrons for both half equations are the same (lowest common multiple). In this case the lowest common multiple is 6, so we need to multiply the reduction half equation by 2, and the oxidation half equation by 3. We can then add the oxidation and reduction half equations together to get an overall redox equation in acidic medium. balance redox reaction overall balanced in acidic medium 2. Balance Redox Equation in Alkaline Medium If we want to balance the redox reaction in alkaline medium, an additional step is required which is to add OH- to neutralise the H+. From the balanced redox equation in acidic medium that we have done previously, there are 8H+ on the left hand side of the equation. Therefore we need to add H- on the left side to neutralise the 8H+, forming 8H2O. We also need to add H- on the right side to keep the equation balanced. Usually there will be H2O on both sides of the equation to cancel out. This will give us the overall balanced redox equation in alkaline medium: balance redox reaction balanced in alkaline medium Watch this video tutorial to learn how to balance a redox equation step-by-step! Topic: Redox Titrations, Physical Chemistry, JC, H2, A Level Chemistry, Singapore You can also view this with screenshots and detailed explanation. Do check out the following for more video lessons: at at If you are looking for H2 Chemistry Tuition, do consider taking up my . /episode/index/show/alevelchemistrylesson/id/7515956

How to draw Amino Acids and Proteins 11/11/2018

Electrophilic Addition Mechanism for Alkenes 11/11/2018

How to compare Basicity of Amines, Phenylamines and Amides 11/11/2018

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