Chemistry I Honors Acids and Bases Lesson #1 Introduction • Welcome to Acids and Bases ! • This unit will use a lot of the concepts covered in the Solutions unit, but also contains some very new material. • We will look at the characteristics and behaviors, and reactions of both types of compounds. First – Some History and Vocabulary • In the mid-1800’s , scientists discovered that some compounds would dissolve in water and the resulting solution would conduct electricity. • Compounds that behaved this way were classified as electrolytes. • Ionic compounds (like salt as NaCl, potassium nitrate as K(NO3), etc) were found to be electrolytes. • It was also taken as “common knowledge” that covalent compounds did not produce solutions that conducted electricity – they were nonelectrolytes. How Does This Occur? • Even though science could identify which compounds were electrolytes, there was no understanding about what happens when one of these compounds dissolves that could account for the electrical conductivity. A Hypothesis • In the late 1800’s, a Ph.D. candidate from Denmark proposed that ionic compounds break apart to form charged particles when dissolved in water. • He named the process “Electrolytic Dissociation”. • The idea was not well received – it was believed at the time that compounds were built from neutral atoms – and the electron had not yet been discovered by Thomson. Eventually… • Sub-atomic particles are discovered, and the theory of Electrolytic Dissociation is accepted. • In the early 1900’s the Nobel Prize in Chemistry is awarded to Svante Arrhenius for his theory of electrolytic dissociation. What This Means • Ionic compounds like NaCl are electrolytes because they dissociate in water to form charged ions. NaCl (s) Na+1 (aq) + Cl-1 (aq) • Covalent compounds like sugar are nonelectrolytes because they are molecular – their dissolving yields complete molecules that do not have a charge. C6H12O6 (s) C6H12O6 (aq) What Does This Have to do with Acids? • Acids are covalent compounds. • We use ionic strategy to write their formulas (SOCCR), but they are covalent compounds (technically polar covalent). • But (and this is huge), they are electrolytes. • Solutions of acids conduct electricity – as if they are ionic – but they aren’t… • And, since they are covalent, they do not dissociate into ions like ionic compounds do. Where do the Ions Come from in Acids? • If acid solutions conduct electricity, something has to occur to produce ions in solution. • Arrhenius proposes (later found correct) that Acids actually react with the water to produce ions. • This reaction is named “Ionization” – the action of forming ions. The Reaction • We have learned that Acids have the general formula “HX” , where the “X” part is either a nonmetal element or a polyatomic ion. • From bonding, the H and the X share a pair of electrons in a covalent bond. • During the Ionization reaction, the H – X bond is broken and the H becomes attached to the two lone pairs on the Oxygen atom of the water molecule. • But, (and this is critical) the X atom keeps both electrons from the H – X bond. Hydronium • The H atom leaves his electron behind (with the X atom), meaning he now has a +1 charge. • The X atom has gained the H’s electron, meaning that he now has an extra electron and as a result now has a -1 charge. • The H+1 particle connects to the water molecule, forming an ion that has the formula H3O+1 • This ion is called the “hydronium ion”. Equation • In equation form, the ionization of an acid HX will look like this: HX + H2O H3O+1 (aq) + X-1 (aq) • All acids that have the chemical format “HX” will react the same way – will all form hydronium and an anion. • Example – Hydrochloric Acid HCl + H2O H3O+1 (aq) + Cl-1 (aq) A Visual More Examples • For Hydrobromic Acid: HBr + H2O H3O+1 (aq) + Br-1 (aq) It works for oxy-acids too! • For Nitric Acid: H(NO3) + H2O H3O+1 (aq) + (NO3)-1 (aq) Some Quick Vocabulary • Monoprotic – this is an acid that only has 1 H atom bonded to the “X” part of the acid molecule. • Examples would be – – – – HCl HBr H(NO3) H(C2H3O2) • Diprotic – these are acids that have 2 H atoms bonded to the “X” part of the acid molecule. • Examples would be – H2(SO4) – H2(CO3) Diprotic Ionizations • Important to remember that during any ionization, the acid loses 1 H atom to a water, forming a single hydronium ion. • This holds true for diprotic acids, but they are able to undergo two separate ionizations. #1) H2(SO4) + H2O H3O+1 + H(SO4)-1 #2) H(SO4)-1 + H2O H3O+1 + (SO4)-2 • Look carefully through the sequence presented above. See if you can write all three ionization equations for phosphoric acid: H3(PO4) What About Bases? • Bases are easier to deal with. • Arrhenius defined bases as “Metal Hydroxides” that dissociate in water to produce a metal cation and the hydroxide anion. • The dissociation is exactly the same as we learned in the Solutions Lesson #1. • There are “monobasic” compounds that have one hydroxide attached to the metal cation. • And there are “dibasic” compounds that have two hydroxides attached to the metal cation. Monobasic and Dibasic Dissociations • Na(OH) is monobasic. (got the formula from SOCCR) • It’s dissociation looks like this: Na(OH) (s) Na+1 (aq) + (OH)-1 (aq) • Mg(OH)2 is dibasic. (again, formula from SOCCR) • It’s dissociation looks like this: Mg(OH)2 (s) Mg+2 (aq) + 2 (OH)-1 (aq) Summary • Acids: – Are electrolytes – Have the formula HX – Ionize in water to form hydronium – Can be monoprotic, diprotic, or triprotic • Bases: – Are electrolytes – Have the formula M(OH) – Dissociate in water to for aqueous ions – most importantly the OH-1 ion. – Can be monobasic or dibasic