Fundamental Inorganic Chemistry PDF

***[Fundamental Inorganic Chemistry ]*** **Reasources:** \- Only for this lecture go to LabPlus ( Davy 7th Floor) **What is inorganic Chemistry:** \- The chemistry of everything thats not carbon-based plus carbon-based moecules less reduced that carbon dioxide Yellow = In all living organisims...

***[Fundamental Inorganic Chemistry ]*** **Reasources:** \- Only for this lecture go to LabPlus ( Davy 7th Floor) **What is inorganic Chemistry:** \- The chemistry of everything thats not carbon-based plus carbon-based moecules less reduced that carbon dioxide Yellow = In all living organisims **Salts + Ions** \- Salts are 2 ions forming something with no charge that will usually be a crystalline solid \- Ions are atoms plus or minus elections to give a net charge \- **Positively** charged ions are called **Cations** - these have lost elections \- **Negatively** charged ions are called **Anions** - These have gained elections \- We talk about the charge as the valency of the ion E.g a 2+ cation is a divalent cation a 3- anion is a trivalent anion **Salt and Compound Names:** \- The first half of the name is the cation, the second half is from the anion: E.g Sodium Chloride, Potassium nitrate \- Some cations are forms of their conjugate bases: \+ Ammonia (NH3) \>\>\> Ammonium (NH4+) \- Sometimes we specify the oxidation state (charge on cation) using Roman numerals in parentheses:\ + Iron (II) Sulphate is FeSO4 aka ferrous sulphate \+ Iron (III) Sulphate is Fe2(SO4) aka ferric sulphate \- Highest oxidation state = -ic Lowest = -ous \- Anion names come from their conjugate acids \+ Nitric acid (HNO3) \>\>\>\> Nitrate (NO3-) \+ Hyrochloric acid (HCI) \>\>\>\> Chloride (CL-) **Acids deprotonate to give atoms:** \- If an acid loses a hydrogen ion we say it has ionised or deprotonated HCI + H2O \>\>\>\> H3O + Cl- It can be written as HCI \>\>\>\> H+ +CI- but you cant have a free H+ ion in solution as they react with water and form H3O+ ions but we conventionally shorthand this to H+ most of the time. \- Some acids are polyprotic meaning they have more than one proton to lose and this gives a range of anions varying the pH. **Oganic Acids Work in the Same Way:** \- Acetic acid (CH3COOH) + H2O \>\>\>\> acetate (CH3COO-) + H3O+ \- We can neautralise the solution with a base (e.g sodium hydroxide, NaOH) to obtain the E.g sodium salt of the acid - NaCH3COO \- This applies to all the organic acids, many of hich are **polyprotic (Acid capable of losing more than a single proton per molecule)** as we have already talked about. E.g Siccinic acid is properly butan-1, 4-dioc acid so has 2 -COOH groups, each of which can deprotonate: \+ HOOCC2H4COOH + H2O \>\>\>\> HOOCC2H4COO- + HH2O+ \+ HOOCC2H4COO- + H20 \>\>\>\> OOCC2H4COO- + HH20 \- The deprotonisation of all acids (organic and inorganic) is pH dependent \- The pH value at which 50% is the acid and 50% is the anion id the pK \- Acctic acid pK = 4.756 - this means you have a 1 M solution at a pH of 4.756, there will be 0.5 M acctic acid and 0.5 M acctate. If the solution was pH 6.756 or above, there would be 1 M acetate and at pH 2.756 or below,1 M acctic acid. In the range in between, there are variable solutions. \- For polyprotic acids there are multiple values - one for each proton. E.g succinic acid has a pKa1 = 4.2 and pKa2 = 5.6 - so at physiological pH (pH7.2), the fully ionised form - succinate - dominates with smaller amount of hydrogensuccinate \- Same idea applies to E.g sulfuric acid (H2SO4) for bisulfate (HSO4-) and sulfate (SO4 2-) ions. \- - - - Acetic acid \-\-\-\-\-- Acetate acid \- - - - Carbon dioxide/Carbonic acid \-\-\-\-\-\-- Bicarbonate \...\...\.... Carbonate **Acids and Bases:** \- Multiple definitions: Arrhenius "acids produce H3O+ ions, bases produce OHions" Brønsted-Lowry "acids donate H+ ions, bases accept H+ ions\" Lewis "acids donate 2ε , bases accept 2ε \- In modern parlance, IUPAC defines \"acid and \"base to mean Arrhenius/ Bronsted-Lowry acids and bases and we only mean Lewis-Theory ones when we explicitly say \" Lewis acid/base\" \- H atom = 1p +1ε H+ ion = 1p ( Cannot physically exist in aqeous solution - H3O+ is formed) We thus colloquially call hydrogen ions protons **Acids:** \- Acetic acid (CH3 COOH) donates protons froming acetate ions: CH3COOH +H20 \>\>\>\> CH3COO- + H30+ can generalise this as AH \>\>\>\> A- + H+ Acetate is the conjugate base of acetic acid **Bases:** \- Ammonia (NH3) accepts protons forming ammonium ions**:** NH3 + H20 \>\>\>\> NH4+ +OH- can generalise this as B + H+ \>\>\>\> BH+ ammonia is the conjugate base of ammonium **Water:** \- Has a neutral pH of 7, therefore it react with its self acting as both an acid and a base. \- H2O + H2O ↔ H3O+ + OH water ↔ hydronium ion + hydroxide ion **-** In acidic solution (pH\7) there are more OH- than H30+ ions pH electrodes dont work at high and low pH due to there not just being H30+ ions **pH:** \- a measure of the H30+ in solution - we oftern shorthand this to H+ and intend that to mean all hydrogen ions like H30+ \- pH + -log10\[H+\] \- \[x\] means molor concentration of x \- we can also use pOH for OH- ions (pOH = -log10\[OH-\] \- Aqueous solutions cannot go above pH 14 but can go to -2 pH **Strong and Weak Acid:** **-** Nothing to do with the concentration of the solution Basd on the pKa and thus a measure of how much of the ion vs free acid is pressent: how easily they dissociate ( deprotonate/ionise) \- Strong acid = pKa \ 0 organic acids are always weak acids hydrochloric acid (pKa = -5.9) - strong mineral acid nitric acid (pKa = -1.4) - strong mineral acid nitrous acid (pKa = +3.2) -- weak mineral acid acetic acid (pKa = +4.8) -- weak organic acid **pH Buffers:** **-** They keep pH in a narrow range around their pKa so even if (H30+) changes, the pH wont actually change \- E.g the phosphate bufferes used in physiological work ar pH 7.2 use the pKa2 of the orthophosphoric acid which covers: H2PO4- + H20 \ HPO4 2- + H30+ \- So if a solution comparason is made 0.5M H2PO4- and 0.5 M HPO4 2-, will be in equilibrium at a pH of 7.2. If something (.g metabolism) causes an increase in H30+ the equilibrium will shift to the left. Therefore HPO4 2- will decrease in concentraion and H2PO4- will increase, mopping up the extra H30+, ensuring pH does not move. Visa Versa