CEM 142 Exam 2

what needs to happen to conduct electricity

what needs to happen to conduct electricity

moving charges

describe chemical reactions and what happens to define it as a chemical reaction

• the chemical formula changes

• at the molecular level

  • some bonds break and some form

  • the overall chemical identity of the substances changes

• it involves the rearrangement of atoms

  • atoms are conserved

  • same # of atoms in the reactants as the products, they’re just rearranged

  • the connections (bonds) between the atoms change

    • break or form

• reactions involve changes in location of valence electrons

• reactions and reactivity can be predicted by understanding how the electrons are distributed in the reactants

• two types:

  • acid-base

  • redox

describe a phase change

• the formula remains the SAME

• IMFs are either formed or broken/overcome

• the formula/element is the same in all phases

• interactions/IMFs are what changes

describe a solution process

• dissolving/making solutions can sometimes be a chemical reaction

• sugar/salt is just a simple solution process

  • same structures and formulas

  • no breaking/forming new ions

    • both states have ions, they just aren’t touching each other when in water

• HCL dissolving is a chemical reaction

  • the products are different formulas than the reactants

• the formation of a solution is sometimes a chemical reaction

important to note about H⁺

• it’s a proton

  • removed an electron

• there’s no such thing as H⁺ by itself in an aqueous solution

• H⁺ is a very small, highly charged species

  • it’s always surrounded by several water molecules (a solvation shell)

• the hydronium ion(H₃O⁺) is a better way to represent H⁺

• if you see H⁺ assume it’s H₃O⁺

describe an Arrhenius acid

• dissolves in water to give H+ ( a proton)

• a bond was broken to get to H+

• since we produced H+, that tells us that we started with an acid

describe an arrhenius base

• dissolves in water to give ^-OH (hydroxide)

• a bond was broken to give -OH

describe the arrhenius acid-base model

• the simplest model

• has limited applications due to its simplicity

• this model ignores the role of water

• anytime an arrhenius acid reacts with an arrhenius base, some kind of salt + water is produced

what is the detailed ionic equation and when would we use it mose?

• shows the details

• shows what ions are present (includes all the molecules)

• shows what species are actually present in solution

• used for arrhenius model mainly

what are spectator ions

• species that don’t change during the chemical reaction

• think of the detailed ionic equation

• since they don’t change, they don’t take part in the reaction

what is the net ionic equation and when do we mainly use it?

• still shows ions, but not all of them

  • just the overall

• shows only the reaction taking place

• used mainly with arrhenius model

formal charge equation reminder

# of valence electrons (on neutral atom) - # of bonds (on atom we’re looking at) - 3 of lone electrons (acutal # of dots, not pairs)

describe bronsted-lowry acids

• a proton (H+) donor

describe bronsted-lowry bases

• a proton (H+) acceptor

what are conjugate acids and bases

• every conjugate acid has a conjugate base and every base has a conjugate acid

• if you’re asking for a conjugate acid, that means the element is a base, basses accept protons

• if you’re asking for a conjugate base that means the element is an acid, acids donate a proton

• “conjugates” since the equilibrium arrow means the reaction can go in either direction

• there will always be conjugate acids/bases since equilibrium and reaction arrows are used interchangeably and technically every reaction is irreversible and therefore in equilibrium, it just might be completely to one side

how does an acid become a conjugate base (bronsted)

the acid donates a proton, what remains of the acid is called the conjugate base

how does a base become a conjugate base (bronsted)

• the base accepts a proton to become the conjugate acid

• what’s left after the base donates a proton is called the conjugate acid

how do acids and bases react (bronsted)? example: HCL and H₂O

• the opposite partial charges of HCL and H2O attract

• the proton that gets transferred is ‘assisted’ by the lone pair on the O

  • the H+ doesn’t just fall off the Cl and then hop onto the water

• in order for H+ to be transferred the bond between HCL must break and a new bond between the H+ and the O forms

• the proton transfers when the molecules collide

  • collisions is what makes this happen

  • molecules collide in a certain way that make it so the partial + (aka H) is attracted to the partial -

  • opposite charges attract which causes the molecules to collide

  • the acid donates a proton to the base

  • one bond breaks and a new one forms

• during collision a proton is transferred from the acid to the base

describe the bronsted-lowry acids and bases model

• more general, a broader model

• a more useful definition

• allows more molecules to be bases, they aren’t limited to hydroxide (-OH)

• good for reactions where protons are transferred and where water is the solvent… but it’s still limited

describe how water can be an acid or a base

• it’s amphiprotic

  • can be both an acid or base

• water can react with water

• it could go in both directions

• if there’s given the same amount of products/reactants, there will be more neutral water molecules than ions in water

  • some ions are present but not that many

  • ions do exist since H2O molecules can react with each other

• evidence: the light board in class

  • the lights didn’t light up so there was no electricity and not enough moving ions to conduct electricity

describe lewis acids

• electron pair acceptors

• have a place to accept the electrons

• this doesn’t have to be a proton (as an acceptor) but can be

  • all bronsted acids are lewis acids

• species with an empty or partially empty orbital (of available energy) are also lewis acids

  • orbitals that can accept electrons

  • group/column 3 and most transition metals

• electron sink- accepts the electrons

describe lewis bases

• electron pair donors

• have available pair of electrons to donate into a bond (formed with the lewis acid)

• to have a lewis base you need to have a lone pair of electrons

• electron source

• initiator of the reaction

describe the lewis acid and base model (example: HCl + H₂O ←→ H₃O⁺ + Cl^-)

• think of acid-base reactions as electron pair donations

• this model encompasses and explains the Bronsted-Lowry and the Arrhenius theory

• BUT we consider the base the electron source and the initiator of the reaction

• this is an even broader model that incorporates reactions where there is no proton transfer (like group 3 with empty p orbitals)

• we use arrow pushing for lewis

• the base donates a lone pair to an H (most times) to form a new bond

example with HCl + H₂O ←→ H₃O⁺ + Cl^-:

• HCl is the acid and H2O is the base

• 2 electrons on the base (water) will be donated to the acid, when this happens a bond will form

  • the base initiates the reaction by donating a pair of electrons

  • the lone pair on the base turns from a lone pair of electrons to a bonding pair of electrons between the O and the new H

• since the H can’t have 2 bonds, it transfers its bonding pair of electrons to the Cl to become a lone pair

  • the acid accepts the electron pair

what’s the difference between bronsted and lewis (specifically in terms of drawings)

bronsted:

• deals with attractions and proton transfer

• in the drawing include dashed line between the partially negative and positive atoms to show an electrostatic attraction

• dashed line between the hydrogen to a lone pair

lewis:

• deals with electron transfer

• lone pair of electrons is donated to the H to form a new bond

• base initiates reaction

• use arrow pushing

• 1st arrow comes from the lone pair on the base and goes to the H on the acid

• 2nd arrow comes from the bonding pair on the H to the acid to make the conjugate

how can BH₃ act as an acid when NH₃ is the base

• base initiates reaction

• B has 2p² hybridization

  • there’s an empty p orbital

  • the empty p orbital allows for 4 bonds, so none will have to break for this reaction

• all of the old bonds remain in tact

• 1st arrow comes from the lone pair on the N and go to the Br (this is since H is similar to C in terms of electronegativity and so the Hs actually have a partial negative charge and the Br has a positive)

• there is no second arrow

• product: the n is bonded to the br as one big molecule and the br has a negative formal charge and the N a positive 1 formal charge

• most transition metals and column 3 have an empty orbital that allow for this

how to spot an acid

• all have Hs attached to an electronegative atom

• acidic Hs are bonded to electronegative elements

  • like O, Cl, Br, I and sometimes N, or F

  • the H will have a partial positive charge since it’s bonded to an electronegative atom

• the H-X bond is weakened by interactions with the solvent

  • once the acid is in a solution, the partial positive H is going to be attracted to the solvent due to the partial charge created from the electronegative atom

  • this attraction is what’ll weaken the bond between X-H

• the conjugate base, X^-, is stable

  • since X is electronegative, it doesn’t mind having the extra lone pair or electrons and doesn’t mind the negative charge, it can stabilize since it’s electronegative

• note: if X is already negative and has an extra set of lone pairs, it doesn’t want another, even if it’s an electronegative atom. it would make the molecule more unstable

comparing acids

acids want to accept electrons/give up H+

• the more electronegative an atom, the bigger the partial charge on the (most times) H bonded to it. the Hs are more attracted to the solvent (usually water) which weakens the H-X (acid) bond and it’ll be easier to break

  • the weakest acids have the strongest conjugate bases and also the least stable conjugate bases

  • the strongest acids have the weakest conjugate bases as they’re the most stable

    • they can better stabilize the negative charge due to a stronger electronegativity

how to spot a base

• all have a lone pair of electrons that are available to form a new bond

• bases HAVE to have a lone pair of electrons

• this lone pair is the lone pair that’s donated in order to form a new bond

• they should be (relatively more) stable when the new bond forms

  • the stronger the base, the weaker the conjugate acid and the more stable the conjugate acid is

• typically, electronegative atoms have lone pairs… but electronegative atoms like to hold onto their electrons

  • so we’re looking for a perfect balance

  • has a lone pair but not holding on too tightly (like N)

• bases donate a pair of electrons to becom a bonding pair

why don’t we include CH₄ or Ne in acid/base comparisons

CH₄

• is not an acid

• no partial charge, the C and H have almost the same electronegativity

Ne

• noble gas, can’t form bonds

• can’t act as a base even though it has all those lone pairs

• also, it’s very electronegative and unwilling to donate electrons to form a bond anyways

why do acid base reactions occur?

• they begin with the partially negative end of one molecule interacting with the partially positive end of the other and they eventually collide

• we can show the reaction mechanism using curved arrows (lewis)

what are electrophiles and nucleophiles an extension of

• of lewis theory

  • type of acid-base ration

• the partial positive isn’t on the Hydrogen… it’s on a CARBON

describe electrophile

• a lewis acid

• electron loving or negative charge loving

• when the partial positive part of a molecule is a Carbon, we call the molecule an electrophile

  • has to be a partial Positive charge so it’ll be attracted to the electrons

• accept electrons

• the positive has to be on the Carbon, the carbon is the electrophilic site

describe nucleophile

• a lewis base

• nucelus loving or positive charge loving

• an electrophile reacts with a nucleophile

• donates electrons

describe CH₃Br (electrophile) reacting with NH₃ (nucleophile)

*disregard light-colored arrow on bottom left

• when the partial positive isn’t on a Hydrogen

• this is called methylation

  • added a methyl group to a N

• CH₃ = methyl group

what is the transition state (use CH₃Br and NH₃ as an example)

• we have both things happening at the same time

• as one bond forms, the other one breaks

• the thing that exists as we transition from the reactants into the products

• this is what the reactants go through as they transition from reactants to products

what does strong mean for acids/bases

• strong acids are fully ionized in a solution

  • every molecule breaks into their separate ions

• strong acids would make a light bulb glow more since there would be more ions to have around as charges

• evidence: using the light with the circuits (in class) and compare brightness

• no equilibrium arrow really since it fully ionizes

what does weak mean for acids/bases

• not fully ionized in solution

  • only some molecules break into ions

• only partially ionize in solution which causes the light bulb to glow relatively dimly at the same concentration

• there is more of the molecular form than the ionized form

• there are still equilibrium arrows but ones longer and ones shorted

  • longer: more of whatever it’s pointing towards

dilute vs concentrated

• only tell about the amount of acid/base, not about the strength

• could have a concentrated/dilute strong/weak acid/base

• strong and weak doesn’t always matter if there’s more solute (aka concentrated) there’s more ions that can move

describe acid strength across a row

• increases across a row (left → right)

• due to electronegativity the electronegativity increasing across a row

• a F-H (in row 2 for example) is the most polarized so the H will have a greater partial positive charge and therefore H is more likely to attract a base by interacting with its lone pair

  • as the acid goes into solution (like H₂O as the base) there’s a STRONG interaction between the acidic H and the O of the water

    • this is what weakens the F-H bond

  • the F^- conjugate base is the most stable due to F’s high electronegativity. The extra pair of electrons is attracted closer to the nucleus and is therefore more stable

    • the conjugate base is stable (it can hold the negative charge) which means that electronegative atoms are most likely to be bonded to acidic Hs

  • the more electronegative the atom, the less likely it’ll donate/share it’s electrons and it’ll be attracted to more electrons

what type of evidence would you use for describing acid strength across a row in terms of HF for an example

• the F is the most electronegative (compared to the other atoms in the row) so H will have the biggest partial positive charge and will interact the most strongly with the solvent/base

OR

• F^- is the most stable conjugate base

  • meaning it’s the weakest base and tied to the strongest acid

describe acid strength down a group and what it depends on

• cannot use electronegativity to explain acid strength down a group

• think thermodynamics

  • strong acids completely ionize (super - ∆G)

  • weak acids partially ionize (kinda - ∆G)

• the stronger the acid (and the more it ionizes) the farther right the reaction will go

• the more thermodynamically favorable a process is, the more - ∆G

acid strength depends on:

• the enthalpy change ∆H when added to water

  • the strength of the H-X bond (that you’re breaking and forming)

  • the stability of the X^- conjugate base

• the entropy change ∆S when added to water

  • the number of arrangements (before and after the reaction) for the ions AND water molecules

• ∆s and ∆H contribute to ∆G

what impacts enthalpy ∆H when comparing acid strength down a row but we don’t focus on? and why don’t we focus on it?

• when an acid (HX) is placed in water and it dissolves and ionizes

when it dissolves:

• it’s a solution process

• IMFs are overcome in the solute and solvent

• solute-solute and solvent-solvent interactions are overcome

  • requiring energy

• solute-solvent interactions form

  • releasing energy

• ∆H from these 2 process is relatively small since IMFs are weak, so we ignore it in our analysis but it still exists/happens

what impacts enthalpy ∆H for acid strength down a column

when HX ionizes

• HX bonds break

  • requires energy (+∆H)

  • not the same for all the acids

    • they all have different bond strengths

• O-H bonds form

  • releases energy (-∆H)

  • when H₂O is the solvent

  • an O-H bond will be a certain strength no matter what the acid is

• ∆H from these 2 processes is relatively large (bonds are stronger than IMFs)

• the energy released upon forming OH bonds with water will be the same no matter what the acid is

• the energy absorbed upon breaking and HX bond will depend on the strength of that bond

bond enthalpy: the energy it takes to break the HX bond in the gas phase (∆H = +)

• the more overlap between orbitals (aka the more similar size orbitals and therefore atoms) the stronger the bond is

• a bigger atom, meaning a smaller % of the orbital is overlapping, the weaker the bond

• this means it requires less energy to break HI than HF (for example) and ∆H is the least positive for HI

• forming an OH bond releases energy and this energy is more than it takes to break the HX bond

  • more energy is released (from the OH bond) than absorbed (from the HX bond) overall so ∆H will be negative for all, it’s a matter of how negative

  • the most -∆H contributes to a more -∆G

• ∆H = OH + HX

  • OH: always the same and very negative since it releases energy

  • depends on acid and always + since it breaks and requires energy

what impacts entropy ∆S for acid strength down a group

• ∆G = ∆H - T∆S

  • subtracting a (-) means adding a number

  • subtracting a (+) means minusing a number

  • a more (+) ∆S will contribute to a more -∆G

• determined by ion sizes

• example comparing F^- and I^-

  • they both have a -1 charge but the (-) chare is more spread out over different-sized atoms

  • the charge is more concentrated in F than in I which causes stronger interactions between F^- and H₂O

  • water can’t have as many arrangements since more are ‘locked’ in around F than I

• smaller ions (like F) attract water molecules tightly which limits the water molecules’ positions (so entropy decreases -∆S)

• large ions don’t attract water molecules as tightly so the water molecules can occupy more positions (so a more +∆S as the entropy increases)

put together the ideas of enthalpy and entropy to describe acid strength down a column

• the more - ∆G the more the reaction lies to the right and the more the molecules break into ions so the stronger the acid

• acid strength increases down a column

  • due to enthalpic and entropic effects

  • cannot use electronegativity

Enthalpic Effect

• the bond strength HX decreases down a group (weaker bonds) due to the difference in size of H and X

• the difference in size between H and X gets bigger down a group

  • there’s less overlap between H and X orbitals when they’re different sizes

• The HX bond is weaker and easier to break

  • takes/requires less energy

Entropic Effect

• entropy decreases when water molecules solvate small, highly charged ions

  • water molecules get locked in place and have fewer arrangements

determine acid strength if H is bonded to the same type of atom in different molecules

reminder:

• the strongest acid has the weakest and most stable conjugate base

• it’s stable because it’s weak and doesn’t need to react

the strongest acid will have resonance structures

• the lone pair on the resonance could become a bonding pair but the C can’t have 5 bonds so one of the other bonds becomes a lone pair on the other Oxygen

  • the negative charge can be considered to be ½ on each oxygen

  • a resonance hybrid

    • this better stabilizes the negative charge since it’s spread out over more atoms

    • which means the conjugate base is the most stable → weaker conjugate base → strongest acid

**induction effect may work for other types of examples too

describe how strong acids have weak conjugate bases

• the stronger the acid, the weaker its conjugate base

• conjugate bases can be stabilized (made weaker) by spreading out the negative charge over several atoms

  • through resonance structures

compare ^-OH to H₂O for base strength

-OH is a stronger base

• the negative on the -OH means it has extra electrons and therefore more willing to donate a pair of electrons (lewis model)

compare NH₃ to H₂O for base strength

NH3 is a stronger base

• O and H are in the same row so you can use electronegativity

• the N is less electronegative than O so it’s not pulling on the lone pair of electrons as hard and is more willing to donate/share the lone paire

  • likewise: the O is more electronegative and is pulling on the lone pair more so it’s less willing to share

which is the stronger base when comparing ^-CH₃ to ^-NH₂ to ^-OH to F^- and why

use lewis structures

• use electronegativity since they’re the same row

• F is the most electronegative and least willing to share its electrons

• C is the least electronegative and most willing to share its electrons so it’s the stronger base

OR conjugate acids

• F is the most electronegative so it gives the biggest partial positive charge on H so it interacts the most strongly with the solvent/base and weakens the FH bond

how to determine the direction of acid-base reactions when at an equilibrium

stronger acid + stronger base ← ———> waker acid + weaker base

• there will always be more of the weaker acid and base

• the reaction can go both ways (hence equilibrium arrows)

  • but it doesn’t go in both directions equally

  • we say that the equilibrium lies to one side or the other

  • more reactants or products are present at equilibrium

• the stronger acid and base are always on one side together and the weaker acid and base are on the other

list of strong acids

assume weak acid for anything else unless specified

*completely ionizes

• HCL

• HBr

• HI

• HNO₃

• H₂SO₄

list of strong bases

*completely ionizes

group 1 or 2 with hydroxide (OH)

• example: NaOH

what’s water’s concentration

55.5 M

how to find % ionization

% ionized = (ionized/total) x 100

• total could be found by ionized + ionized or by how much acid/base you started with/were given

how to find pH and how to find [H₃O⁺]

pH = -log[H₃O⁺]

• log base 10

• note: the pH can go beyond 1-14 at both ends

[H₃O⁺] = 10^-pH

what is pH

• tells the concentration of H⁺ (but really H₃O⁺) in solution

  • this is since H+ doesn’t exist by itself in the solution

• pH tells how acidic or basic a solution is

  • high pH: more basic

    • less [H₃O⁺]

    • more [^-OH]

  • low pH: more acidic

    • more [H₃O⁺]

    • less [^-OH]

what is auto-ionization of water and what does this tell us about the [H₃O⁺] and the [^-OH]

• when water reacts with itself

• this is rapid and reversible in water:

  • there are more neutral water molecules than ions

  • evidence: light board in class

• in pure water the concentration of H₃O⁺ equals the concentration of ^-OH

• [ ] = concentration (M)

• at 25˚C the concentration of each ion is 1 × 10^-7M

  • [H₃O⁺] = [^-OH] = 1 × 10^-7M

how to find [H₃O⁺] [^-OH] from each other at 25˚C

[H₃O⁺] x [^-OH] = 1 × 10^-14

• at 25˚C

[H₃O⁺] = 1 × 10^-14 / [^-OH]

[^-OH] = 1 × 10^-14 / [H₃O⁺]

what does pH depend on

TEMPERATURE

this solution (in the pic) is neutral:

• the concentration of H₃O⁺ and -OH are equal so the solution is neutral

• the solution is pure water

• pH depends on temperature

acidic vs basic vs neutral

• deals with relative concentrations of H3O+ and -OH

• acid = more H3O+

• base = more -OH

[H₃O⁺] and [^-OH] relationship

how to find [H₃O⁺] when given a strong acid or how to find [^-OH] when given a strong base

*the same holds true from the image for strong bases

• strong bases fully ionize

  • not all strong bases have a super high pH

    • likewise with acids at low pHs

  • pH depends on [H₃O⁺] and [^-OH], not necessarily strong vs weak

describe redox reactions and oxidation vs reduction

oxidation-reduction reactions

• both oxidation and reduction are happening in the reaction

  • one thing is being oxidized and another is reduced

oxidation

• if an atom loses one or more electrons it’s being oxidized

• lose valence electrons

reduction

• gaining 1 or more electrons

• reducing in charge

*the electrons from what’s being oxidized are transported/going to what’s being reduced

OLI RIG

• Oxidation Is Loss (of electrons)

• Reduction Is Gain (of electrons)

what are oxidation state/numbers

• used to keep track of where electrons are before and after reaction

• do not confuse with formal charges

  • both keep track of electrons but tell you different things

• they’re a way of electron book-keeping

  • allows you to see whether the element has lost or gained electrons (or a bigger share of the electrons) during the reaction

• if oxidation #s change during a reaction then it’s a redox reaction

  • if they don’t, then it’s not a redox

common rules for oxidation numbers

1. atoms in pure elements = 0

   1. examples: H₂, O₂, Na

2. ions = ion charge

   1. example: NaCl

   2. Na = +1

   3. Cl = -1

   4. total/sum = 0

3. oxygen = -2

4. hydrogen = 1

5. the sum of a molecule’s oxidation #’s should equal 0

   1. unless there’s a formal charge on the molecule, then the sum of the oxidation #’s should equal the charge

trick for assigning oxidation #s

• withdraw all of the electrons in each bond to the most electronegative atom

  • remember that this doesn’t actually happen

  • once you pull the electrons to the more electronegative atom you calculate formal charge (to find the oxidation number)

formal charge vs. oxidation numbers

**both are ways of electron book-keeping

formal charge

• treating bonding electrons like one electron belongs to one atom and one electron belongs to the other

• used to draw the most stable lewis structures

  • better to have fewer and lower # formal charges in lewis structures

oxidation number

• treating the bonding electrons like they both belong to the more electronegative atom

• used to determine if a reaction is a redox reaction or not

• keeps track of electrons and where they’re transferred in a redox reaction

  • we ‘give’ all the electrons to the most electronegative atom in a bond

• they decide if a redox reaction happens or not

what are resonance structures

• if you can take 1 structure and use curved arrows to get to the other structure, that’s how you know you have resonance

• since the charge is spread out/split between atoms, it can cause the acid to be stronger since it stabilizes the negative charge better on the conjugate base, making it a weaker conjugate base

  • also, the spread out negative causes a weaker conjugate base since the lower magnitude of charge won’t attract a proton as easily

  • a resonance hybrid

    • this better stabilizes the negative charge since it’s spread out over more atoms

    • which means the conjugate base is the most stable → weaker conjugate base → strongest acid

why do we use multiple models of acid-base reactions?

many models help us predict and explain outcomes. generally, we use the simplest model that’ll work in a given solution. even though all acid-base reactions can be described using the lewis model, sometimes it is simpler and easier to use the bronsted-lowry model. this is partially true for reactions in aqueous solution where proton transfer is the easiest way to think about the reaction

2 ways of explaining stronger acids using the bronsted model

1. discuss electronegativity and how H has a bigger partial positive so it’s more likely to be given away and is more attracted to the base/solvent

2. electronegativity: after it donates the proton, it can better stabilize the negative charge

   1. the stronger the acid, the weaker the conjugate base and more stable the conjugate base

3. thermodynamics (down a column for acids)

bronsted to explain stronger bases

• bases have to accept protons

• the atom has room for another bond (aka a lone pair) and the atom has a partial negative charge which will attract the proton

lewis model to explain stronger acids

• acid accepts a lone pair from the base

• the more electronegative atom will have stronger interactions with its lone pair so it’s less likely to give up its electrons

lewis model to explain stronger bases

• bases are electron pair donors

• the less electronegative atom is more likely to give up its electron pairs

describe the beginning, middle, and end of reactions

beginning

• acid-base are interacting (usually through hydrogen bonding interactions)

middle/transition state

• the bond forming and breaking happens at the same time

• the HX (acid) bond breaks as the HY (base) bond forms

  • the H doesn’t fall off and reattach

• the H+ bonded to X is attracted to the lone pair on Y and the HX bond is weakened until it’s overcome

end

• they’re ionized

what happens to the H3O+ if we add solute to a solution and the pH stays the same

• the H3O+ stays the same too

• the solute didn’t act as an acid or a base

  • the solute doesn’t react with the water/solvent

strong vs weak acid/base using concentrations

• if the given is strong (acid or base) then the H3O+ or -OH concentration will equal the amount of the given (acid or base)

• then you can solve for pH

• the pH is what the pH should be if it were a strong acid or base

  • you can compare when the concentration of the substances are the same

Why is CN- not an acid?

• it’s already negative

• more electrons would make it unstable

• wants to get rid of those extra electrons

what are inductive effects

• when electronegative atoms in a molecule pull the other electrons (electron density) toward themselves

• for a base (for example) to be more stable, the charge on the base is better stabilized since it’s more spread out

• the electronegative atoms pull the electron density toward themselves making the base less likely to share its electrons

• the less spread out the charge = the more basic